| Topic: |
Science > Physics |
| User: |
"HC" |
| Date: |
06 Dec 2007 01:51:16 AM |
| Object: |
Electrolysis, newbie, redux |
Hey, all. Thank you for your time to read this. I'm sorry for posting
this to .physics and .chem but I'm not sure which is really the better
forum for this type of question.
So, a long time ago (in this same galaxy, no far away at all) I was
monkeying with electrolysis of water to produce hydrogen and oxygen.
I had some problems (chlorine gas instead of O2) and posted about them
and then worked on it a little more and then let it sit for about 2
years because I was frustrated. Well, after a long hiatus I'm back
after this. I am trying to correct some of my earlier mistakes (I was
using NaCL for my electrolyte) and make some progress. I purchased
some gouging rods (carbon rods used with arc welders to remove metal
material) as electrodes and am now using baking soda instead of NaCL.
The gouging rods are copper clad. I fired the thing up using tap
water (from a well here in north central Texas) that has some NaCL
(from the water softener) and, at about 13 volts I got some nice
bubbles from the cathode (H2, if I'm not a complete idiot), and some
small bubbles at the anode (O2, maybe...but since I made chlorine last
time...confidence is in short supply). I did seem to get some opaque,
milky something at the anode. I had a meter in-line and was getting
about 0.10 amps across the solution. I shut it down and dissolved
about 1/2 tablespoon of baking soda in to the water and fired it back
up. The current went to almost 0.40 amps and the milky discharge from
the anode continued. After a very short time it seemed to have a
bluish hue, so I think it's some kind of copper-derivative (from the
copper cladding on the electrodes). The cathode didn't seem to give
any such discharge but was giving off gas vigorously, presumably
hydrogen. I quit the experiment, if you can call something this
rudimentary that, and the water is slightly bluish (after about 10
minutes of run time).
Does anybody have any general suggestions on what is going on here?
What I would like is to have a solution of something in water that,
when electricity is passed through it, will form hydrogen and oxygen
and will leave a solution behind that is short a little water that is
easily replenished with distilled water. I am, and probably always
will be, an amateur and a tinkerer, so please forgive my ineptitude.
Thank you.
--HC
.
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 08:16:26 PM |
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On Dec 6, 1:51 am, HC <hboo...@gte.net> wrote:
Hey, all. Thank you for your time to read this. I'm sorry for posting
this to .physics and .chem but I'm not sure which is really the better
forum for this type of question.
So, a long time ago (in this same galaxy, no far away at all) I was
monkeying with electrolysis of water to produce hydrogen and oxygen.
I had some problems (chlorine gas instead of O2) and posted about them
and then worked on it a little more and then let it sit for about 2
years because I was frustrated. Well, after a long hiatus I'm back
after this. I am trying to correct some of my earlier mistakes (I was
using NaCL for my electrolyte) and make some progress. I purchased
some gouging rods (carbon rods used with arc welders to remove metal
material) as electrodes and am now using baking soda instead of NaCL.
The gouging rods are copper clad. I fired the thing up using tap
water (from a well here in north central Texas) that has some NaCL
(from the water softener) and, at about 13 volts I got some nice
bubbles from the cathode (H2, if I'm not a complete idiot), and some
small bubbles at the anode (O2, maybe...but since I made chlorine last
time...confidence is in short supply). I did seem to get some opaque,
milky something at the anode. I had a meter in-line and was getting
about 0.10 amps across the solution. I shut it down and dissolved
about 1/2 tablespoon of baking soda in to the water and fired it back
up. The current went to almost 0.40 amps and the milky discharge from
the anode continued. After a very short time it seemed to have a
bluish hue, so I think it's some kind of copper-derivative (from the
copper cladding on the electrodes). The cathode didn't seem to give
any such discharge but was giving off gas vigorously, presumably
hydrogen. I quit the experiment, if you can call something this
rudimentary that, and the water is slightly bluish (after about 10
minutes of run time).
Does anybody have any general suggestions on what is going on here?
What I would like is to have a solution of something in water that,
when electricity is passed through it, will form hydrogen and oxygen
and will leave a solution behind that is short a little water that is
easily replenished with distilled water. I am, and probably always
will be, an amateur and a tinkerer, so please forgive my ineptitude.
Thank you.
--HC
An update on what I've done so far (maybe this will help somebody):
First, I stopped using the proto-board as my power supply. It seemed
it was the choke point. Moving to a regular wall-wart yields
consistent results and current through the solution. The wall-wart
I'm now using is a 120 VAC to 6 VDC, 1.66 amp adapter (I have no idea
what this came from and the 1.66 amps is quite high in my experience
for a wall-wart). It works quite well, doesn't overheat, and works
consistently. I'm using plastic gallon water jugs with two holes
pierced in the top for the electrodes to pass through. I fill the jug
about half-way with tap water. My tap water is from a well with a
softener so it does have some
Moving on, to the solutions I've tried. Two years ago I tried table
salt (NaCl) which worked but gave me H2 and Cl2 gases instead of H2
and O2. It's been suggested that continuing to use the solution for a
long enough period of time would finally rid the solution of all the
Cl it had to offer, leaving NaOH in solution which would be a good
electrolyte for electrolysis of water. I have not tried that and
concern about the health risks of Cl gas make me not want to run it
that long. I could attempt to find some NaOH to put into a virgin
solution but I'm not working that angle currently.
In the last week or so of messing with this again I have tried several
solutions, all with carbon rods for electrodes. The carbon rods are
gouging rods used with arc welders and are copper clad and about
5/16th of an inch in diameter.
Recent attempt 1: Baking soda in water with copper clad electrodes. I
was using the proto-board at about 13 volts. I got about 1+ amp of
current across this but it was choking (current would drop quite a bit
after a little time (few minutes). Seemed the proto-board PS was
shutting down (probably some kind of internal protection). I switched
to the wall-wart (see above). That helped but the copper cladding was
failing (turning blue-green and contaminating the water and leaving
some kind of solids in the bottom of the jug).
Recent attempt 2: Baking soda in water but with the copper cladding
removed from the gouging rods (carefully peeled off the copper
cladding with a knife and patience). I cannot say that this had any
effect on the resultant gases (I'm not testing either electrode for
its gas as they're in the same plastic jug and IF they're H2 and O2
they'd be in perfect explosive ratio so I'm avoiding sparks and flints
and such) but it DID keep the water cleaner. The gas from the
positive terminal did appear in larger bubbles than the foamy/opaque
stuff when the copper cladding was on. However, the water did turn a
grayish color. I'm not sure why that was as I wouldn't think the
carbon of the rod would dissolve in the water, much less actually
become a mixture of uniform color.
Recent attempt 3: Vinegar and bare carbon rods. The vinegar was too
diluted to begin with, I think, and by the time I added small amounts
to the tap water it was a lost cause; too little reaction, too little
gas, too little current. :-/
Recent attempt 4: Sulfuric Acid (H2SO4) with bare carbon rods (skinned
another pair of the rods). This worked very well as gas was given off
from both electrodes, no discoloration of the water occurred, and,
with sufficient acid in the mixture, the current was as high as was
seen with the baking soda. Specifically, I mixed acid that was
intended to fill a lead-acid battery (but had never been inside one)
with tap water. More precisely, I administered (sounds kinda like
something Sherlock Holmes would say) 250 drops of the acid into
approximately 1/2 gallon of tap water. Yes, I counted the drops, ten
at a time. I noted how many sets of 10 I added to the water and
watched as the current increased from 0.04 amps to about 0.52 amps
from the wall-wart. I could add more acid, I'm sure, but I'm not sure
how much more conductive the water would get, and for now it satisfies
my need to know if sulfuric acid will work. I like its results the
best of the batch. Oh, I used a drinking straw and my finger as the
dropper, FWIW, to measure the acid into the water.
I attempted to get Glauber's Salt (sodium sulfate) from my local Wal-
mart this evening. That was good for a few smirks. The woman working
the pharmacy, and presumably the pharmacist on duty from her garb,
didn't have a clue what Glauber's Salt was. She suggested, when I
asked her where I might find some, that I try in the spice
section. :-/ So, after a few more attempts I gave up and looked for
myself in the laxatives section but with no luck. Next I'll try
online or I my skip it and go for sodium bisulfate.
Anyway, that's it. I hope this will help anybody who comes along and
is looking for information as I was when I started this post.
Thanks to all who have aided with information. It is appreciated.
--HC
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| User: "N:dlzc D:aol T:com \dlzc" |
|
| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 09:20:35 PM |
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|
Dear HC:
"HC" <hboothe@gte.net> wrote in message
news:eb2ff333-6863-4afe-a8f6-954cd7c2ebc2@q77g2000hsh.googlegroups.com...
....
I attempted to get Glauber's Salt (sodium sulfate)
from my local Wal-mart this evening. That was
good for a few smirks. The woman working
the pharmacy, and presumably the pharmacist on
duty from her garb, didn't have a clue what
Glauber's Salt was.
Try epsom's salt... magnesium sulfate.
David A. Smith
.
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
12 Dec 2007 11:32:29 AM |
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On Dec 11, 9:20 pm, "N:dlzc D:aol T:com \(dlzc\)" <dl...@cox.net>
wrote:
Dear HC:
"HC" <hboo...@gte.net> wrote in message
news:eb2ff333-6863-4afe-a8f6-954cd7c2ebc2@q77g2000hsh.googlegroups.com...
...
I attempted to get Glauber's Salt (sodium sulfate)
from my local Wal-mart this evening. That was
good for a few smirks. The woman working
the pharmacy, and presumably the pharmacist on
duty from her garb, didn't have a clue what
Glauber's Salt was.
Try epsom's salt... magnesium sulfate.
David A. Smith
David, would that be more effective than the Glauber's salt or sodium
bisulfate (the pH Minus, IIRC)?
--HC
.
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| User: "N:dlzc D:aol T:com \dlzc" |
|
| Title: Re: Electrolysis, newbie, redux |
12 Dec 2007 06:17:38 PM |
|
|
Dear HC:
"HC" <hboothe@gte.net> wrote in message
news:61a61548-ea42-44eb-8850-f569178fb44a@i12g2000prf.googlegroups.com...
On Dec 11, 9:20 pm, "N:dlzc D:aol T:com \(dlzc\)"
<dl...@cox.net>
wrote:
Dear HC:
"HC" <hboo...@gte.net> wrote in message
news:eb2ff333-6863-4afe-a8f6-954cd7c2ebc2@q77g2000hsh.googlegroups.com...
...
I attempted to get Glauber's Salt (sodium sulfate)
from my local Wal-mart this evening. That was
good for a few smirks. The woman working
the pharmacy, and presumably the pharmacist on
duty from her garb, didn't have a clue what
Glauber's Salt was.
Try epsom's salt... magnesium sulfate.
David, would that be more effective than the
Glauber's salt or sodium bisulfate (the pH Minus,
IIRC)?
Try it. The druggist / pharmacist will at least not look at you
like you were daft.
David A. Smith
.
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
23 Dec 2007 07:33:06 PM |
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On Dec 6, 1:51 am, HC <hboo...@gte.net> wrote:
Hey, all. Thank you for your time to read this. I'm sorry for posting
this to .physics and .chem but I'm not sure which is really the better
forum for this type of question.
So, a long time ago (in this same galaxy, no far away at all) I was
monkeying with electrolysis of water to produce hydrogen and oxygen.
I had some problems (chlorine gas instead of O2) and posted about them
and then worked on it a little more and then let it sit for about 2
years because I was frustrated. Well, after a long hiatus I'm back
after this. I am trying to correct some of my earlier mistakes (I was
using NaCL for my electrolyte) and make some progress. I purchased
some gouging rods (carbon rods used with arc welders to remove metal
material) as electrodes and am now using baking soda instead of NaCL.
The gouging rods are copper clad. I fired the thing up using tap
water (from a well here in north central Texas) that has some NaCL
(from the water softener) and, at about 13 volts I got some nice
bubbles from the cathode (H2, if I'm not a complete idiot), and some
small bubbles at the anode (O2, maybe...but since I made chlorine last
time...confidence is in short supply). I did seem to get some opaque,
milky something at the anode. I had a meter in-line and was getting
about 0.10 amps across the solution. I shut it down and dissolved
about 1/2 tablespoon of baking soda in to the water and fired it back
up. The current went to almost 0.40 amps and the milky discharge from
the anode continued. After a very short time it seemed to have a
bluish hue, so I think it's some kind of copper-derivative (from the
copper cladding on the electrodes). The cathode didn't seem to give
any such discharge but was giving off gas vigorously, presumably
hydrogen. I quit the experiment, if you can call something this
rudimentary that, and the water is slightly bluish (after about 10
minutes of run time).
Does anybody have any general suggestions on what is going on here?
What I would like is to have a solution of something in water that,
when electricity is passed through it, will form hydrogen and oxygen
and will leave a solution behind that is short a little water that is
easily replenished with distilled water. I am, and probably always
will be, an amateur and a tinkerer, so please forgive my ineptitude.
Thank you.
--HC
Another update.
Yesterday I finally tried some Sodium Bisulfate. I got it at Home
Depot in their outdoor lawn and garden area (it's for pools so it's
off season) as pH Minus. It was about 10 dollars for 7 pounds. I put
3 spoonfuls (disposable plastic dinnerware spoons) into about 1/2
gallon of water (each time I do these experiments I use the same
amount of water to try to have some consistency between different
tries). Using the same wall-wart I've been using I got 0.66 amps
through the solution and what seemed like more gas than I've been
getting with other chemicals. Of course, presumably, I could get
better conductivity in the solution of sulfuric acid if I added more
of it to the test container I have for the acid but since I was able
to add such a small amount of the pH Minus to the water and get better
results I'll probably stick with it. That may not be clear: I have
multiple containers I am using. Each container has only one
electrolyte in it. I have one for sulfuric acid, one for sodium
bisulfate (and others for other chemicals I've tried).
In any case, I was reading some information someone posted online
about creating an electrolysis apparatus for use in cars and they were
using stainless steel wall plates from Home Depot as the electrodes.
I have no intention of using the electrolysis apparatus in a car but
the concept of using stainless steel plates sounds good so I grabbed a
couple today and will try them after Christmas. We'll see if that
works well and if it doesn't corrode.
What I've also read about that sounds interesting is that some people
are using PWM (pulse-width modulators) to process the power going to
the electrolysis apparatuses. Does anybody know if that would make a
difference on the efficiency of an electrolysis rig? I would *think*
that the PWM would just provide down times in the voltage thereby
lowering the amount of gas produced and NOT that it would somehow
increase the production of gas over the same voltage/current being
applied constantly to the same solution. Is that correct or is there
some phenomena that cause the use of intermittent (pulsed) DC to
produce more gas than constant DC? Before I go to the trouble of
obtaining a power MOSFET to monkey with this I'd like some
information.
Thank you all, again, for your help and time and information.
Merry Christmas.
--HC
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| User: "Bill Penrose" |
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| Title: Re: Electrolysis, newbie, redux |
23 Dec 2007 08:24:27 PM |
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On Dec 23, 6:33 pm, HC <hboo...@gte.net> wrote:
In any case, I was reading some information someone posted online
about creating an electrolysis apparatus for use in cars and they were
using stainless steel wall plates from Home Depot as the electrodes.
I've used stainless steel electrodes before with good luck. Be sure
there's not much chloride around, or the steel can corrode. Stainless
steel is not entirely stainless.
DB
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 09:15:13 PM |
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On Dec 6, 1:51 am, HC <hboo...@gte.net> wrote:
Hey, all. Thank you for your time to read this. I'm sorry for posting
this to .physics and .chem but I'm not sure which is really the better
forum for this type of question.
So, a long time ago (in this same galaxy, no far away at all) I was
monkeying with electrolysis of water to produce hydrogen and oxygen.
I had some problems (chlorine gas instead of O2) and posted about them
and then worked on it a little more and then let it sit for about 2
years because I was frustrated. Well, after a long hiatus I'm back
after this. I am trying to correct some of my earlier mistakes (I was
using NaCL for my electrolyte) and make some progress. I purchased
some gouging rods (carbon rods used with arc welders to remove metal
material) as electrodes and am now using baking soda instead of NaCL.
The gouging rods are copper clad. I fired the thing up using tap
water (from a well here in north central Texas) that has some NaCL
(from the water softener) and, at about 13 volts I got some nice
bubbles from the cathode (H2, if I'm not a complete idiot), and some
small bubbles at the anode (O2, maybe...but since I made chlorine last
time...confidence is in short supply). I did seem to get some opaque,
milky something at the anode. I had a meter in-line and was getting
about 0.10 amps across the solution. I shut it down and dissolved
about 1/2 tablespoon of baking soda in to the water and fired it back
up. The current went to almost 0.40 amps and the milky discharge from
the anode continued. After a very short time it seemed to have a
bluish hue, so I think it's some kind of copper-derivative (from the
copper cladding on the electrodes). The cathode didn't seem to give
any such discharge but was giving off gas vigorously, presumably
hydrogen. I quit the experiment, if you can call something this
rudimentary that, and the water is slightly bluish (after about 10
minutes of run time).
Does anybody have any general suggestions on what is going on here?
What I would like is to have a solution of something in water that,
when electricity is passed through it, will form hydrogen and oxygen
and will leave a solution behind that is short a little water that is
easily replenished with distilled water. I am, and probably always
will be, an amateur and a tinkerer, so please forgive my ineptitude.
Thank you.
--HC
More experience. I stripped the copper cladding off a couple more
gouging rods so I have bare carbon rods. I ran them in to another
container of water with baking soda. I added about 2 tablespoons to a
half gallon of water. I applied about 13 volts to the solution and
got about 1.5 amps across the solution. I got vigorous bubbles at
both electrodes for a while but then the current dropped to about 0.4
amps and the bubbles subsided quite a bit. Any suggestions or
knowledge about what might be occurring here? I stopped the
electrolysis for a few moments and then started it again. When I
restarted it the current jumped back up to the 1.5 amp range but
quickly returned to the lower 0.4 amp range. I tried vinegar but that
was not any good (too slow and too little gas produced). Unless
somebody says something to the contrary I will try some sulphuric
acid.]
--HC
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| User: "N:dlzc D:aol T:com \dlzc" |
|
| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 09:32:02 PM |
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|
Dear HC:
"HC" <hboothe@gte.net> wrote in message
news:ab3dce98-7651-4f4c-9a2d-1e226118a282@e67g2000hsc.googlegroups.com...
....
Unless somebody says something to the contrary I will
try some sulphuric acid.
Something.
What happens if you remove all the oxygen from H2SO4? Better try
this one in a well ventillated area, and your hand on the kill
switch.
David A. Smith
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 11:28:52 PM |
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On Dec 6, 9:32 pm, "N:dlzc D:aol T:com \(dlzc\)" <dl...@cox.net>
wrote:
Dear HC:
"HC" <hboo...@gte.net> wrote in message
news:ab3dce98-7651-4f4c-9a2d-1e226118a282@e67g2000hsc.googlegroups.com...
...
Unless somebody says something to the contrary I will
try some sulphuric acid.
Something.
What happens if you remove all the oxygen from H2SO4? Better try
this one in a well ventillated area, and your hand on the kill
switch.
David A. Smith
David, thank you for your reply. If I'm understanding your post,
there is a point at which the chemical reaction between H20, H2SO4,
and e will cease to produce H2 and 2O2 because some component of the
sulphuric acid (H2SO4) will be depleted. At that point some other
chemical(s) will be given off. Is that correct? If that is correct
then the solution of water and sulphuric acid has a finite period of
time (specifically it has a finite amount of H2 and 2O2 to give off)
before it quits working in that way and begins to work in another
(that is, it starts to give off another chemical(s)). If what I am
saying/understanding is correct then sulphuric acid would not meet the
requirements that I desire of a solution from which I can "cook"
hydrogen and oxygen and replenish by simply adding more distilled
water.
Is what I've said correct/accurate? What would the solution, once the
component(s) of the sulphuric acid have been depleted and the reaction
changes, give off?
Again, thank you for your response.
--HC
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 11:43:48 PM |
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On Dec 6, 11:28 pm, HC <hboo...@gte.net> wrote:
On Dec 6, 9:32 pm, "N:dlzc D:aol T:com \(dlzc\)" <dl...@cox.net>
wrote:
Dear HC:
"HC" <hboo...@gte.net> wrote in message
news:ab3dce98-7651-4f4c-9a2d-1e226118a282@e67g2000hsc.googlegroups.com...
...
Unless somebody says something to the contrary I will
try some sulphuric acid.
Something.
What happens if you remove all the oxygen from H2SO4? Better try
this one in a well ventillated area, and your hand on the kill
switch.
David A. Smith
David, thank you for your reply. If I'm understanding your post,
there is a point at which the chemical reaction between H20, H2SO4,
and e will cease to produce H2 and 2O2 because some component of the
sulphuric acid (H2SO4) will be depleted. At that point some other
chemical(s) will be given off. Is that correct? If that is correct
then the solution of water and sulphuric acid has a finite period of
time (specifically it has a finite amount of H2 and 2O2 to give off)
before it quits working in that way and begins to work in another
(that is, it starts to give off another chemical(s)). If what I am
saying/understanding is correct then sulphuric acid would not meet the
requirements that I desire of a solution from which I can "cook"
hydrogen and oxygen and replenish by simply adding more distilled
water.
Is what I've said correct/accurate? What would the solution, once the
component(s) of the sulphuric acid have been depleted and the reaction
changes, give off?
Again, thank you for your response.
--HC
Oops, sorry, NOT H2 and 2O2 but rather 2H2 and O2.
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| User: "John Savage" |
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| Title: Re: Electrolysis, newbie, redux |
10 Dec 2007 01:38:08 AM |
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HC <hboothe@gte.net> writes:
both electrodes for a while but then the current dropped to about 0.4
amps and the bubbles subsided quite a bit. Any suggestions or
knowledge about what might be occurring here? I stopped the
electrolysis for a few moments and then started it again. When I
restarted it the current jumped back up to the 1.5 amp range but
quickly returned to the lower 0.4 amp range. I tried vinegar but that
Bicarb soda is as good as any electrolyte. The limiting factor is your
type of electrode. Carbon is good initially because it has a large
surface area so it has a large area in contact with the liquid. But
after a few minutes of operation carbon can be seen to be a poor choice
for exactly the same reason: tiny bubbles of liberated gas adhere to the
carbon rod---and a carbon rod cloaked in a film of non-conductive gas
does not make for good electrical conduction! This can only be solved
by incorporating some method for continually dislodging the gas from
the electrodes, either by a stirrer or a shaker, and using electrodes
of large surface area. Also, that milky stream of H2 gas is akin to a
submerged foam, a mix of tiny bubbles and liquid, and it too contributes
to reducing the conducting properties of the electrolyte . I haven't
heard a good explanation accounting for the big difference in 0 and H
bubbles.
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
--
John Savage (my news address is not valid for email)
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| User: "John Savage" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 07:20:34 PM |
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John Savage <rookswood@suburbian.com.au> writes:
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
Sorry, what I wrote there was confusing. I meant to say you definitely
can't use Al for both electrodes. [Then I went on to recount an expt I
did using ac from 0.1 Hz up to 50 Hz to determine at what frequency an
electrolytic rectifier (constructed using a pair of aluminium electrodes
in bicarb sode solution) ceased any rectifying behaviour and became just
a resistance. My recollection is that it was somewhere around 5-10Hz,
certainly way less than the mains frequency here; the implication being
that it is not feasible to employ an electrolytic rectifier on mains ac.]
--
John Savage (my news address is not valid for email)
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 08:31:12 PM |
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On Dec 11, 7:20 pm, John Savage <rooksw...@suburbian.com.au> wrote:
John Savage <rooksw...@suburbian.com.au> writes:
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
Sorry, what I wrote there was confusing. I meant to say you definitely
can't use Al for both electrodes. [Then I went on to recount an expt I
did using ac from 0.1 Hz up to 50 Hz to determine at what frequency an
electrolytic rectifier (constructed using a pair of aluminium electrodes
in bicarb sode solution) ceased any rectifying behaviour and became just
a resistance. My recollection is that it was somewhere around 5-10Hz,
certainly way less than the mains frequency here; the implication being
that it is not feasible to employ an electrolytic rectifier on mains ac.]
--
John Savage (my news address is not valid for email)
Hey, John, that's an interesting idea. I'm not sure of the
intricacies involved but it sounds like a cool concept. Sorry it
didn't work out.
--HC
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
10 Dec 2007 06:33:59 PM |
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On Dec 10, 1:38 am, John Savage <rooksw...@suburbian.com.au> wrote:
HC <hboo...@gte.net> writes:
both electrodes for a while but then the current dropped to about 0.4
amps and the bubbles subsided quite a bit. Any suggestions or
knowledge about what might be occurring here? I stopped the
electrolysis for a few moments and then started it again. When I
restarted it the current jumped back up to the 1.5 amp range but
quickly returned to the lower 0.4 amp range. I tried vinegar but that
Bicarb soda is as good as any electrolyte. The limiting factor is your
type of electrode. Carbon is good initially because it has a large
surface area so it has a large area in contact with the liquid. But
after a few minutes of operation carbon can be seen to be a poor choice
for exactly the same reason: tiny bubbles of liberated gas adhere to the
carbon rod---and a carbon rod cloaked in a film of non-conductive gas
does not make for good electrical conduction! This can only be solved
by incorporating some method for continually dislodging the gas from
the electrodes, either by a stirrer or a shaker, and using electrodes
of large surface area. Also, that milky stream of H2 gas is akin to a
submerged foam, a mix of tiny bubbles and liquid, and it too contributes
to reducing the conducting properties of the electrolyte . I haven't
heard a good explanation accounting for the big difference in 0 and H
bubbles.
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
--
John Savage (my news address is not valid for email)
Hey, John, thank you for your reply and time. That's a very
interesting observation; I had not thought of that (about the film of
non-conducting gas). So, if I understand correctly what you're
saying, I could use bicarb soda (baking soda) just as effectively (or
pretty close) as any other electrolyte for what I'm doing? Instead of
using carbon rods (gouging rods) I could use led sheets which would
not tend to become insulated with gas or coverd in oxidization
(aluminum).
If I use lead I would assume I would not want to try the H2SO4 that
has been mentioned and that I've been planning on trying (I obtained
some today that is new and was intended to fill a lead-acid battery
but not tried it yet). I assume that because lead and acid react with
on another. Correct? Also, would the lead contaminate the water in
any appreciable way? I mean, would this contaminate the water so much
that it would be hazard to health or the environment? I'm not a big
tree hugger but I would like to know more about what negative outcome
there might be.
Thank you again for the information.
--HC
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| User: "John Savage" |
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| Title: Re: Electrolysis, newbie, redux |
14 Dec 2007 06:31:30 AM |
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HC <hboothe@gte.net> writes:
Instead of
using carbon rods (gouging rods) I could use led sheets which would
not tend to become insulated with gas or coverd in oxidization
(aluminum).
Whatever you use it will still become cloaked in bubbles, especially
when you seem to be aiming for a higher volume of gas production. But
for d-i-y it is much easier to form a large electrode surface out
of metal than carbon, and bubbles seem to adhere less to metal than
lampblack rods. The ultimate might be lead sheets dipping into an
electrolyte contained in old aluminium saucepan so the saucepan sides
and base comprise one of the electrodes. Or use a soda can that you
have scrubbed inside with something abrasive to remove the plastic
film from the aluminium.
If I use lead I would assume I would not want to try the H2SO4 that
has been mentioned and that I've been planning on trying (I obtained
some today that is new and was intended to fill a lead-acid battery
but not tried it yet). I assume that because lead and acid react with
on another. Correct?
No, I think you'll find lead doesn't react with dilute H2SO4. Use two
lead electrodes and you'll in fact be charging a battery. The surface
of one electrode will stay grey (spongy lead), the other will turn
chocolate (lead oxide). But both these coatings are conductive, so the
current will keep flowing and the cell gassing. After you're finished,
connect a low-voltage globe between the two and it should light for a
few seconds until the cell goes flat! The difference here is that with
dil H2SO4 you are liberating lots of gas, while the proper charging of
a car battery should not generate much gas. (The holes in the caps are
not big enough to let a lot of gas escape! Some car batteries are
sealed and don't gas.)
Also, would the lead contaminate the water in any appreciable way?
No, negligible I'd say. But the electrodes WILL lose metal through
erosion, not corrosion. If you use carbon electrodes, after an hour
or so you will see the carbon surface has become rough, some of it
having been eroded by agitation by the bubbles. The same will happen
for metal, but to a lesser extent. The eroded carbon shows up as gunk
in the cell. Think about it: within tiny depressions in the carbon a
molecule of water suddenly explodes one-thousand or so times in volume
into a mole of H2 gas. The pressure wave breaks off tiny fragments of
the rod, eventually leaving it visibly pitted. (The process sounds like
cavitation, or the reverse of it.)
As an experiment, you could try adding a *tiny* (say, one tenth of
a drop) of detergent to the cell, to reduce the surface tension of the
water without it being sufficient to form suds! This may encourage the
bubbles to leave the electrodes faster because that is the limiting
condition to how fast you can liberate gases for a particular setup.
There may be some household product better than detergent, a surfactant
that is not sudsing, but I can't think of one.
--
John Savage (my news address is not valid for email)
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| User: "N:dlzc D:aol T:com \dlzc" |
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| Title: Re: Electrolysis, newbie, redux |
14 Dec 2007 07:06:22 AM |
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Dear John Savage:
"John Savage" <rookswood@suburbian.com.au> wrote in message
news:071214000113203.14Dec07$rookswood@suburbian.com...
....
(The holes in the caps are not big enough to let a
lot of gas escape! Some car batteries are sealed
and don't gas.)
The size of the hole has nothing to do with whether or not you
are "overcharging" the battery, driving the reaction to liberate
hydrogen. The size of the hole is to minimize evaporation of the
water, so that consumers do not add tap water as fill... as
often.
During normal operation with no cells shorted, the battery is
only ever charged to the point that there is only a tiny bit of
lead sulfate left. No significant hydrogen production will have
occured at this point.
David A. Smith
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| User: "Jim Black" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 01:18:04 PM |
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On Dec 10, 4:33 pm, HC <hboo...@gte.net> wrote:
If I use lead I would assume I would not want to try the H2SO4 that
has been mentioned and that I've been planning on trying (I obtained
some today that is new and was intended to fill a lead-acid battery
but not tried it yet). I assume that because lead and acid react with
on another. Correct? Also, would the lead contaminate the water in
any appreciable way? I mean, would this contaminate the water so much
that it would be hazard to health or the environment? I'm not a big
tree hugger but I would like to know more about what negative outcome
there might be.
Thank you again for the information.
--HC
If you use lead for the positive terminal, you will probably produce
lead compounds in the electrolysis reaction. Whether this would
present a hazard I don't know.
--
Jim E. Black
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| User: "Jim Black" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 01:05:39 PM |
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On Dec 9, 11:38 pm, John Savage <rooksw...@suburbian.com.au> wrote:
HC <hboo...@gte.net> writes:
both electrodes for a while but then the current dropped to about 0.4
amps and the bubbles subsided quite a bit. Any suggestions or
knowledge about what might be occurring here? I stopped the
electrolysis for a few moments and then started it again. When I
restarted it the current jumped back up to the 1.5 amp range but
quickly returned to the lower 0.4 amp range. I tried vinegar but that
Bicarb soda is as good as any electrolyte. The limiting factor is your
type of electrode. Carbon is good initially because it has a large
surface area so it has a large area in contact with the liquid. But
after a few minutes of operation carbon can be seen to be a poor choice
for exactly the same reason: tiny bubbles of liberated gas adhere to the
carbon rod---and a carbon rod cloaked in a film of non-conductive gas
does not make for good electrical conduction! This can only be solved
by incorporating some method for continually dislodging the gas from
the electrodes, either by a stirrer or a shaker, and using electrodes
of large surface area. Also, that milky stream of H2 gas is akin to a
submerged foam, a mix of tiny bubbles and liquid, and it too contributes
to reducing the conducting properties of the electrolyte . I haven't
heard a good explanation accounting for the big difference in 0 and H
bubbles.
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
--
John Savage (my news address is not valid for email)
If you use metal at the positive electrode, then some of the electrons
being removed from that side will come not from a reaction like
2H2O -> O2 + 4H+ + 4e-
but from the metal atoms as they dissolve in the water:
M -> M+ + e-
or
M -> M++ + 2e-
or
M -> M+++ + 3e-
etc., depending on what metal M is. This is the reason for the
reduced oxygen when using metal electrodes.
--
Jim E. Black
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| User: "Salmon Egg" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 08:19:46 PM |
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On 12/11/07 11:05 AM, in article
f718021e-3d18-4781-9642-4813132d700c@e25g2000prg.googlegroups.com, "Jim
Black" <tramspap@yahoo.com> wrote:
If you use metal at the positive electrode, then some of the electrons
being removed from that side will come not from a reaction like
2H2O -> O2 + 4H+ + 4e-
but from the metal atoms as they dissolve in the water:
M -> M+ + e-
or
M -> M++ + 2e-
or
M -> M+++ + 3e-
etc., depending on what metal M is. This is the reason for the
reduced oxygen when using metal electrodes.
The key to what is happening at an electrode is not whether it is positive
or negative but whether it is a cathode or anode. The Daniell or gravity
cell illustrates this. Look up the gravity cell variant in Wickipedia. It is
not clear from the illustration there, but there are copper sulfate crystals
on the bottom of the jar.
To use as a cell, a connection is made between the two electrodes through
the load. Electrons are left behind on the zinc electrode as the zinc
dissolves by turning into soluble zinc ions. These electrons pass through
the external wire and load to the copper electrode. These electrons
neutralize the soluble positive copper ions in solution. These now neutral
atoms get deposited onto the copper electrode. As the copper gets deposited,
negative sulfate ions are left behind. Copper sulfate dissolves replenishing
the copper that was plated out and adding more sulfate ion. This increased
concentration diffuses through the copper sulfate solution, then through the
zinc sulfate solution to the zinc electrode where the sulfate give up their
electrons to the zinc as the zinc dissolves. The zinc ion neutralizes the
additional sulfate ion showing up at the zinc electrode.
When such a cell delivering power, the zinc is getting oxidized. By
definition, that is an anode. There is reduction of copper ion at the copper
electrode. By definition, that is the cathode.
If the circuit is opened, charge builds up on the electrodes to where
sulfate ion are repelled at the zinc an attracted at the copper. In this
case there is no oxidation or reduction at either electrode nd there is no
anode or cathode.
Now suppose that an external dc source in excess of 1.1V is connected to the
cell--positive to the copper and negative to the zinc. In the absence of
side reactions, zinc would get plated out and copper would dissolve. Now the
negative zinc becomes a cathode while the positive copper becomes the anode.
In both cases, the fundamental reactions at the anode and cathode are the
same, irrespective of polarity.
I did not mean to ramble on and on, but once started I had trouble stopping.
Bill
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
11 Dec 2007 08:27:08 PM |
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On Dec 11, 1:05 pm, Jim Black <trams...@yahoo.com> wrote:
On Dec 9, 11:38 pm, John Savage <rooksw...@suburbian.com.au> wrote:
HC <hboo...@gte.net> writes:
both electrodes for a while but then the current dropped to about 0.4
amps and the bubbles subsided quite a bit. Any suggestions or
knowledge about what might be occurring here? I stopped the
electrolysis for a few moments and then started it again. When I
restarted it the current jumped back up to the 1.5 amp range but
quickly returned to the lower 0.4 amp range. I tried vinegar but that
Bicarb soda is as good as any electrolyte. The limiting factor is your
type of electrode. Carbon is good initially because it has a large
surface area so it has a large area in contact with the liquid. But
after a few minutes of operation carbon can be seen to be a poor choice
for exactly the same reason: tiny bubbles of liberated gas adhere to the
carbon rod---and a carbon rod cloaked in a film of non-conductive gas
does not make for good electrical conduction! This can only be solved
by incorporating some method for continually dislodging the gas from
the electrodes, either by a stirrer or a shaker, and using electrodes
of large surface area. Also, that milky stream of H2 gas is akin to a
submerged foam, a mix of tiny bubbles and liquid, and it too contributes
to reducing the conducting properties of the electrolyte . I haven't
heard a good explanation accounting for the big difference in 0 and H
bubbles.
Instead of carbon, you can use sheet lead for both electrodes; or
sheet aluminium for the (-) electrode and lead for the (+). You
possibly can't use Al for both, as at low frequency ac Al reacts with
the O2 bubbles to form an insulating anodising layer of Al oxide. (Yes,
I have experimented with this myself, but just forget what the results
were.)
--
John Savage (my news address is not valid for email)
If you use metal at the positive electrode, then some of the electrons
being removed from that side will come not from a reaction like
2H2O -> O2 + 4H+ + 4e-
but from the metal atoms as they dissolve in the water:
M -> M+ + e-
or
M -> M++ + 2e-
or
M -> M+++ + 3e-
etc., depending on what metal M is. This is the reason for the
reduced oxygen when using metal electrodes.
--
Jim E. Black
Hey, Jim, thanks for your reply. What you've mentioned there reminds
me of some distant information from some chemistry I took long ago;
that is, in most reactions maybe 1 or 2 or 3 electrons might get
swapped around or exist as surplus in anions, but not 4 or more (at
least not usually). I could be wrong, it's been 15 years or so since
I had a chemistry class, but something like that. I remember 1-3
being common, 4 or more not.
Anyway, regardless of my poor recollection of such stuff, what you're
saying makes sense and would account for how electroplating could/does
work. I think John's point was that using carbon as the electrodes
was bad only from the standpoint of how the gas, once formed at the
surface of the electrode, was retained by the surface of the
electrode; that it would form an insulating layer of bubbles. With
the information you're offering here, could I maybe at least improve
the process by replacing the negative electrode with lead or aluminum
which might not have the surface retention of gas like the carbon and
still not force metal ions into the solution?
Thanks again.
--HC
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| User: "Bill Penrose" |
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| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 09:49:55 PM |
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On Dec 6, 7:15 pm, HC <hboo...@gte.net> wrote:
.. Unless
somebody says something to the contrary I will try some sulphuric
acid.]
That's one of the better electrolytes, if dilute enough. Dilute the
acid to about 0.1 N with distilled water. That's 1 mL of concentrated
acid to each 360 mL of water, added dropwise with lots of mixing.
Slowing of current can be due to polarizing of the electrode with gas.
You can overcome this by increasing the voltage, up to a point.
Dangerous Bill
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
06 Dec 2007 11:42:37 PM |
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On Dec 6, 9:49 pm, Bill Penrose <penr...@iit.edu> wrote:
On Dec 6, 7:15 pm, HC <hboo...@gte.net> wrote:
.. Unless
somebody says something to the contrary I will try some sulphuric
acid.]
That's one of the better electrolytes, if dilute enough. Dilute the
acid to about 0.1 N with distilled water. That's 1 mL of concentrated
acid to each 360 mL of water, added dropwise with lots of mixing.
Slowing of current can be due to polarizing of the electrode with gas.
You can overcome this by increasing the voltage, up to a point.
Dangerous Bill
Hey, Bill, thank you for your reply. In the post that appears just
prior to yours there is an allusion to some hazard of using sulphuric
acid. It seems there is a point at which the chemical reaction might
be depleted of a certain component and at which time the reaction will
change with dire consequences. Do you have any information on that?
I continued thinking about the problem I was seeing and thought I
might try a different power supply (I was using a proto board for
electronics as the power supply originally). Fearing there might be
some overload-protection circuitry (that might shut it down if it got
hot) in the device (or God-knows-what to keep monkeys from burning up
the proto board) I tried the SAME solution with a wall-wart that is
labeled as being capable of outputting 6 volts DC at 1.66 amps. I ran
it through an ammeter and came up with only about 0.36 amps, BUT it
ran at that level for over 30 minutes (much longer than the other
power supply could run at 1.5 amps or so). So, either the fancy proto
board has some kind of thermal/current/whatever cutout that allows it
to run at a certain level for only so long and then cuts it back down
OR whatever performance barrier I ran up against only evidences itself
above a certain amount of current/energy and 0.36 amps is insufficient
to bring about that change.
At this point I'm making two gases with electrolysis using baking soda
as the electrolyte. I assume one to be hydrogen and one to be oxygen.
Having removed the copper-cladding from the electrodes I am NOT
getting any green coloring of the water, so I ASSUME I'm getting a
desired reaction (H2 and O2) from the solution. I would still like to
try sulphuric acid, though, but I'm not going to try it until I learn
more about the consequences.
Thank you again for your reply.
--HC
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| User: "Salmon Egg" |
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| Title: Re: Electrolysis, newbie, redux |
07 Dec 2007 05:20:55 PM |
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On 12/6/07 9:42 PM, in article
d1cc4d34-99f7-46e2-ae2d-c7d6f1ab20ec@i12g2000prf.googlegroups.com, "HC"
<hboothe@gte.net> wrote:
At this point I'm making two gases with electrolysis using baking soda
as the electrolyte. I assume one to be hydrogen and one to be oxygen.
Having removed the copper-cladding from the electrodes I am NOT
getting any green coloring of the water, so I ASSUME I'm getting a
desired reaction (H2 and O2) from the solution. I would still like to
try sulphuric acid, though, but I'm not going to try it until I learn
more about the consequences.
You can try using sodium bisulfate instead of the bicarbonate. That will be
safer than sulfuric acid and greatly increase the ion mobility of the
solution.
Sulfuric acid is not too bad if you treat it with respect. You might be able
to get battery acid at an auto supply store. You can get concentrated acid
at a hardware store as a drain cleaner. That, however, really does require
great respect, especially if you do not know what you are doing.
Bill
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
07 Dec 2007 06:23:05 PM |
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On Dec 7, 5:20 pm, Salmon Egg <Salmon...@sbcglobal.net> wrote:
On 12/6/07 9:42 PM, in article
d1cc4d34-99f7-46e2-ae2d-c7d6f1ab2...@i12g2000prf.googlegroups.com, "HC"
<hboo...@gte.net> wrote:
At this point I'm making two gases with electrolysis using baking soda
as the electrolyte. I assume one to be hydrogen and one to be oxygen.
Having removed the copper-cladding from the electrodes I am NOT
getting any green coloring of the water, so I ASSUME I'm getting a
desired reaction (H2 and O2) from the solution. I would still like to
try sulphuric acid, though, but I'm not going to try it until I learn
more about the consequences.
You can try using sodium bisulfate instead of the bicarbonate. That will be
safer than sulfuric acid and greatly increase the ion mobility of the
solution.
Sulfuric acid is not too bad if you treat it with respect. You might be able
to get battery acid at an auto supply store. You can get concentrated acid
at a hardware store as a drain cleaner. That, however, really does require
great respect, especially if you do not know what you are doing.
Bill
Hi, Bill, thank you for your response. I think I'm okay with the acid
(I have handled it successfully in the past with motorcycle batteries
and old car batteries). I would love to have a solution that was
safer AND more effective/efficient so if the sodium bisulfate is
better then maybe I should use that. I read about it a little on
www.wikipedia.org (um, how did humanity survive without deja groups
and wikipedia?). It seems I might be able to find this stuff as a pH
modifier for swimming pools. Is that correct? If not, or even if so,
where would you recommend I get this stuff in a format/purity that
would be acceptable for electrolysis of water?
I am still interested in using the acid, of course, because 1) it's
highly available and 2) the danger seems less important with the
dilution I've seen mentioned in this thread (i.e., very dilute). But,
if sodium bisulfate could be had easily, would be safer, and more
effective/efficient to boot, then I'll go with that.
Thank you again for your response.
--HC
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| User: "Jim Black" |
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| Title: Re: Electrolysis, newbie, redux |
08 Dec 2007 12:55:20 AM |
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On Fri, 7 Dec 2007 16:23:05 -0800 (PST), HC wrote:
On Dec 7, 5:20 pm, Salmon Egg <Salmon...@sbcglobal.net> wrote:
You can try using sodium bisulfate instead of the bicarbonate. That will be
safer than sulfuric acid and greatly increase the ion mobility of the
solution.
Sulfuric acid is not too bad if you treat it with respect. You might be able
to get battery acid at an auto supply store. You can get concentrated acid
at a hardware store as a drain cleaner. That, however, really does require
great respect, especially if you do not know what you are doing.
Bill
Hi, Bill, thank you for your response. I think I'm okay with the acid
(I have handled it successfully in the past with motorcycle batteries
and old car batteries). I would love to have a solution that was
safer AND more effective/efficient so if the sodium bisulfate is
better then maybe I should use that. I read about it a little on
www.wikipedia.org (um, how did humanity survive without deja groups
and wikipedia?). It seems I might be able to find this stuff as a pH
modifier for swimming pools. Is that correct? If not, or even if so,
where would you recommend I get this stuff in a format/purity that
would be acceptable for electrolysis of water?
I've seen the stuff in Wal-Mart under the brand name "hth pH Minus":
http://www.qualityinflatables.com/images/hth61310large.jpg
The "inert ingredients" are sodium sulfate and water:
http://msds.walmartstores.com/cache/39441_1.pdf
--
Jim E. Black (domain in headers)
How to filter out stupid arguments in 40tude Dialog:
!markread,ignore From "Name" +"<email address>"
[X] Watch/Ignore works on subthreads
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| User: "Salmon Egg" |
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| Title: Re: Electrolysis, newbie, redux |
08 Dec 2007 08:31:34 AM |
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On 12/7/07 10:55 PM, in article
1w6xd76zumase$.1uxyc3wxdo4sj$.dlg@40tude.net, "Jim Black"
<fmlast3@organization.edu> wrote:
Hi, Bill, thank you for your response. I think I'm okay with the acid
(I have handled it successfully in the past with motorcycle batteries
and old car batteries). I would love to have a solution that was
safer AND more effective/efficient so if the sodium bisulfate is
better then maybe I should use that. I read about it a little on
www.wikipedia.org (um, how did humanity survive without deja groups
and wikipedia?). It seems I might be able to find this stuff as a pH
modifier for swimming pools. Is that correct? If not, or even if so,
where would you recommend I get this stuff in a format/purity that
would be acceptable for electrolysis of water?
I've seen the stuff in Wal-Mart under the brand name "hth pH Minus":
I sometimes need to adjust the pH of my hyrdroponic nutrients. I use diluted
sulfuric acid derived from drain cleaner. Various government agencies make
it very difficult to obtain what used to be commonly available chemicals.
Usually, products labeled pH up or pH down, command greatly inflated prices
compared to the ingredients used. For hydroponics, you would not want to use
sodium bisulfate because sodium is usually detrimental to plant growth. Some
plants are extremely sensitive. For those applications, potassium bisulfate,
ikf available, would be a better choice.
Bill
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| User: "hanson" |
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| Title: Re: Electrolysis, newbie, redux |
08 Dec 2007 10:36:52 AM |
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"Salmon Egg" <SalmonEgg@sbcglobal.net> wrote in message
news:C37FEAC6.F95D%SalmonEgg@sbcglobal.net...
Various government agencies make it very difficult to
obtain what used to be commonly available chemicals.
Bill
[hanson]
Ain't that wonderfully safe!... AHAHAHA... ahahaha...
So, well then you folks, just keep on rooting for more
environmental protection, for more safety in anything
from food to toys, for more public safety to shield you
from drunk drivers, goons & terrorists and CERTAINLY,
in time, a government agency you will assign a personal
BIG BROTHER to you to protect you who will tell you
what you can & cannot do, have or buy... ahaha...
while you live in the bliss of oblivion, you conveniently
forgetting that it was YOU folks who invited all that...
ahahahaha... AHAHAHAHA...
So, why is it that are you crying "Various government
agencies make it very difficult to obtain what used to be
commonly available chemicals"? ???... All those rights
that were taken away from you is ultimately your own
fault!... ....ahahahaha... ahahaha.. AHAHAHA...
Now watch BigBro's next assignment to protect you &
***** you..."They", whom YOU elected, are already
hatching the next future safety plan for you, by them
sunning their fat asses on Bali, with the great California
enviro *****, Sen. Boxer, heading the parade, to see
to it that YOU WILL PAY for your right to exhale CO2...
ahahaha... She and them will get your money so that you
can exhale a sigh of relieve... Ain't that wonderfully safe!
ahahaha... Here's more, you hopeless mooches... ahaha:
< http://groups.google.com/group/sci.environment/msg/14968cc3ee9939d4 >
Thanks for the laughs... ahahaha... ahahanson
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| User: "V-for-Vendicar" |
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| Title: Re: Electrolysis, newbie, redux |
14 Dec 2007 07:47:02 PM |
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"hanson" <hanson@quick.net> wrote Absolutely nothing.
Meanwhile as the Globe continues to warm...
11 Hottest Years Occurred in Past 13
------------------------------------
- Dave Mosher
LiveScience.com
Thu Dec 13, 4:50 PM ET
New climate data indicate this year might be one of 11 of the hottest years
on
record, all of which have occurred in the past 13 years.
Announced today at United Nations global climate change talks in Bali,
Indonesia, the conclusion stems from global climate records dating back to
1850
and new data collected from January through November 2007. So far, 2007 has
been
the seventh hottest year on record.
The last few days have provided an important platform for debate and
confirms
the need for swift action to combat further rises in global temperatures
because
of human behavior," said Vicky Pope, head of climate predictions at the Met
Office Hadley Center for Climate Change in England, from the Bali
conference.
Scientists and politicians at the conference are discussing plans to reduce
greenhouse gas emissions, which have been linked to rising global
temperatures.
While 2007 is shaping up for seventh hottest in the worldwide category,
different regions have seen more marked increases in temperature. The half
of
the planet north of the equator, for example, may be just shy of an all-time
annual heat record.
2007 was warmer in the Northern Hemisphere, where the year ranks second
warmest,
than the Southern Hemisphere, where it ranks ninth warmest," said Phil
Jones, a
climatologist at the University of East Anglia in England. The year began
with a
weak El Nino -- the warmer relation of La Nina -- and global temperatures
well
above the long-term average. However, since the end of April the La Nina
event
has taken some of the heat out of what could have been an even warmer year.
Across the United Kingdom, 2007 is on course to break all previous warm
temperature records.
Even if December's average temperature is 1.8 degrees Fahrenheit (1 degree
Celsius) below the 1971-2000 long-term average, it would still be the area's
third warmest since 1914. Of the UK's 94 years of local records, 2002
through
2007 are set to become the six warmest years ever for the region.
The top 10 hottest years globally, based on average temperatures, include:
1998 - 32.94 degrees Fahrenheit (0.52 degrees Celsius)
2005 - 32.86 degrees Fahrenheit (0.48 degrees Celsius)
2003 - 32.83 degrees Fahrenheit (0.46 degrees Celsius)
2002 - 32.83 degrees Fahrenheit (0.46 degrees Celsius)
2004 - 32.77 degrees Fahrenheit (0.43 degrees Celsius)
2006 - 32.76 degrees Fahrenheit (0.42 degrees Celsius)
2007 - 32.74 degrees Fahrenheit (0.41 degrees Celsius)
2001 - 32.72 degrees Fahrenheit (0.40 degrees Celsius)
1997 - 32.65 degrees Fahrenheit (0.36 degrees Celsius)
1995 - 32.5 degrees Fahrenheit (0.28 degrees Celsius)
....
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| User: "HC" |
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| Title: Re: Electrolysis, newbie, redux |
08 Dec 2007 03:04:38 AM |
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On Dec 8, 12:55 am, Jim Black <fmla...@organization.edu> wrote:
On Fri, 7 Dec 2007 16:23:05 -0800 (PST), HC wrote:
On Dec 7, 5:20 pm, Salmon Egg <Salmon...@sbcglobal.net> wrote:
You can try using sodium bisulfate instead of the bicarbonate. That will be
safer than sulfuric acid and greatly increase the ion mobility of the
solution.
Sulfuric acid is not too bad if you treat it with respect. You might be able
to get battery acid at an auto supply store. You can get concentrated acid
at a hardware store as a drain cleaner. That, however, really does require
great respect, especially if you do not know what you are doing.
Bill
Hi, Bill, thank you for your response. I think I'm okay with the acid
(I have handled it successfully in the past with motorcycle batteries
and old car batteries). I would love to have a solution that was
safer AND more effective/efficient so if the sodium bisulfate is
better then maybe I should use that. I read about it a little on
www.wikipedia.org(um, how did humanity survive without deja groups
and wikipedia?). It seems I might be able to find this stuff as a pH
modifier for swimming pools. Is that correct? If not, or even if so,
where would you recommend I get this stuff in a format/purity that
would be acceptable for electrolysis of water?
I've seen the stuff in Wal-Mart under the brand name "hth pH Minus":
http://www.qualityinflatables.com/images/hth61310large.jpg
The "inert ingredients" are sodium sulfate and water:
http://msds.walmartstores.com/cache/39441_1.pdf
--
Jim E. Black (domain in headers)
How to filter out stupid arguments in 40tude Dialog:
!markread,ignore From "Name" +"<email address>"
[X] Watch/Ignore works on subthreads- Hide quoted text -
- Show quoted text -
Jim, thank you for your reply and time. That helps a lot. I will
check Wally World for this.
--HC
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| User: "Bill Penrose" |
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| Title: Re: Electrolysis, newbie, redux |
07 Dec 2007 01:43:46 AM |
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On Dec 6, 9:42 pm, HC <hboo...@gte.net> wrote:
... It seems there is a point at which the chemical reaction might
be depleted of a certain component and at which time the reaction will
change with dire consequences. Do you have any information on that?
I can't imagine what that might be. Sulfuric acid is used all the time
for this reaction. If you use concentrated sulfuric acid, you might
get some ozone produced, which is pretty hazardous. The acid should be
quite dilute. Sodium sulfate will also work.
Just make sure the power supply is pure DC, which your proto-board
supply is. You shouldn't be able to harm the power supply. Too much
current drain will pull the voltage down, but those boards are made
for people who make mistakes, so nothing should blow out.
Wall-wart transformers have the problem that they do burn out if their
specs are exceeded. There's no protection. Some surplus providers eg
www.mpja.com will sell you a powerful, protected power supply for a
few dollars.
Remember the classic tests for the gases. The smoldering wood splinter
will burst into flame in the oxygen (at the positive electrode) and
the gas itself will ignite with a pop if it's hydrogen. There should
be no odor.
Dangerous Bill
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