Sciencemadness Discussion Board - Hydrogen from Aluminum and Water Economically

Hydrogen from Aluminum and Water Economically

symboom - 19-9-2011 at 00:10

figuring out info on preforming the experiment

first i propose to start with aluminum and infusing gallium metal into the crystal lattice so the aluminum will react with water.

not sure what the ratio is for 1 pound of aluminum how much gallium is need to cause the protective layer on the aluminum to be disrupted so all the aluminum will react.
generating hydrogen gas

2 now the reaction of aluminum and water forms aluminum hydroxide
"does the gallium react along with aluminum" being more surface area.
and separation: boiling it out cant work gallium metal 3999 °F unlike mercury 674.11 °F but i think liquid gallium should pass right through a filter right?? any suggestions for this?decant off solid from liquid?

then comes the recycling part using molten calcium chloride at 1422F and at that temperature decomposing the aluminum hydroxide to aluminum oxide with loss of water. as aluminum oxide is formed it dissolves in the molten calcium chloride

the + and - electrodes made of carbon argon gas should shield the reaction as the aluminum is molten at 1220F and the CaCl2 is at 1422F
to prevent the aluminum from oxidizing

[Edited on 19-9-2011 by symboom]

blogfast25 - 19-9-2011 at 05:18

You want to sacrifice one of the most expensive metals (gallium) to make hydrogen 'economically'???

Listen son, get some clean Al scraps or some Zn cladding. Add dilute cheap drains unblocker (H2SO4) or cheap patio brick cleaner (HCl) and think of an apparatus to dry and clean the generated hydrogen. Presto!

You even get to keep the resulting salts, for FREE!

[Edited on 19-9-2011 by blogfast25]

Endimion17 - 19-9-2011 at 05:44

He probably read that old media scam article about a guy/team making a car that "runs on water". There was only one detail - it required aluminium-gallium alloy that reduces the hydrogen from water to its elemental state.

Forget gallium, aluminium is far more important here. Just consider the enthalpies. Making freaking aluminium to use it for ... omg I feel so ineffective right now.

blogfast25 - 19-9-2011 at 06:07

Endi:

You're relatively new here but I'm sure you've figured out that we have 'broad church' here, including people that believe uranium rocks are all over the place, one who thought beaches are made of rutile (some are: they're called mines!), one filling up a latex balloon with hydrogen from Al + NaOH and other phantasmagorical schemes. Synboom has previouses, by the way...

bfesser - 19-9-2011 at 08:24

symboom, try doing a little research before posting nonsense.
Ziebarth's doctoral <a href="http://docs.lib.purdue.edu/dissertations/AAI3417965/">thesis</a> on the subject.

bquirky - 19-9-2011 at 09:52

use aluminium as a anode in an electrolysis cell

the bubbling mechanically removes the oxide layer allowing allowing more aluminium to oxidize, eventually you have a beaker full of a white gell mess and more hydrogen than is possible from electrolysis alone.

use it to bug a thermodynamisist. I hate those guys

symboom - 19-9-2011 at 12:32

Quote: Originally posted by blogfast25

You want to sacrifice one of the most expensive metals (gallium) to make hydrogen 'economically'???

Listen son, get some clean Al scraps or some Zn cladding. Add dilute cheap drains unblocker (H2SO4) or cheap patio brick cleaner (HCl) and think of an apparatus to dry and clean the generated hydrogen. Presto!

You even get to keep the resulting salts, for FREE!

[Edited on 19-9-2011 by blogfast25]

i said to recycle the gallium to alloy more aluminum
and the reaction is with water not acid i dont want to deal with the acid. mostly want to try it for the recycling.

and recycling is the most important cause yes galllium is expensive

and the ffc cambridge process for processing metals from their oxides and no it wont be 100% efficient i know this .

separating the gallium metal from aluminum hydroxide from the dead reaction mixture? galium metal can't be heated because it oxidizes in air and gallium oxide cant be reduced in the ffc cambridge process either.
only pure aluminum oxide

[Edited on 19-9-2011 by symboom]

blogfast25 - 19-9-2011 at 12:50

Quote: Originally posted by symboom

Blahdiblahdiblah. Except... that you'll never do this of course!

Quote: Originally posted by symboom

then comes the recycling part using molten calcium chloride at 1422F and at that temperature decomposing the aluminum hydroxide to aluminum oxide with loss of water. as aluminum oxide is formed it dissolves in the molten calcium chloride

the + and - electrodes made of carbon argon gas should shield the reaction as the aluminum is molten at 1220F and the CaCl2 is at 1422F
to prevent the aluminum from oxidizing

[Edited on 19-9-2011 by symboom]

Apart from the almost complete non-feasibility for a hobbyist, do you have any proof that alumina dissolves in molten CaCl2? Al2O3 dissolves in molten Na3AlF6 (cryolite) with a low eutectic and to some extent also in molten CaF2 (fluorite). Don't you think the current Hall-Herould Process would use CaCl2 as a solvent if it was suitable?

Even as a flux for molten Al, CaCl2 isn't very suitable because it has quite high volatility.

You don't have the foggiest idea of what you're talking about, do you?

symboom - 19-9-2011 at 12:57

do you have any proof that alumina dissolves in molten CaCl2?
and yes i do http://www.metalysis.com/
http://www.metalysis.com/products.php
shows what elements can be made from their metal oxide

Endimion17 - 19-9-2011 at 13:24

Wow, this place is awesome. I feel like I'm in elementary school again.

IrC - 19-9-2011 at 16:24

No one considered that reacting the Gallium-Aluminum leaves blobs of shiny liquid Gallium at the bottom of the warm water making it easy to collect? I have done this many times. While warm and wet you can squeeze the leftover residue put into a folded piece of chamois and the Ga comes out just the way Hg does. Try it sometime.

symboom - 19-9-2011 at 19:45

oh thanks that was exactly what i wanted to know so the gallium liquid passes right through the cloth like water? so vacuum filtering the solution of gallium and aluminum hydroxide should clean the gallium right up. and thus separating the gallium metal to be used again and the aluminum hydroxide.

blogfast25 - 20-9-2011 at 06:06

Quote: Originally posted by symboom

do you have any proof that alumina dissolves in molten CaCl2?
and yes i do http://www.metalysis.com/
http://www.metalysis.com/products.php
shows what elements can be made from their metal oxide

Sure. And a fantasist like you is going to reproduce the FCC in his backyard (LOL). To produce 'economical hydrogen' no less. The next Edison you ain't, that much is clear.

Quote: Originally posted by symboom

oh thanks that was exactly what i wanted to know so the gallium liquid passes right through the cloth like water? so vacuum filtering the solution of gallium and aluminum hydroxide should clean the gallium right up. and thus separating the gallium metal to be used again and the aluminum hydroxide.

Gallium is more reactive than aluminium in many ways. The idea your gallium stays behind unaffected is a complete nonsense.

YAAaawwnn...

[Edited on 20-9-2011 by blogfast25]

symboom - 21-9-2011 at 13:43

the process is really not that hard ive already recycled the gallium and as soon as i get my Fresnel lens ill use that to melt the calcium chloride melts at 1421F, and my welder has enough argon to last at least long enough to electrolyze the small sample of aluminum hydroxide i have.
2700F hottest temp at full focal point im sure the lens concentrates sun light hot enough.

ps i know this process is not 100% efficient that is impossible
and hydrogen is only a conveyer of energy it takes more electricity to make hydrogen than it is to burn that if it were to be used to generate electricity again.
and btw hydrogen from electrolysis of water is what i use for if i need hydrogen.
im more doing the experiment because im interested in extraction of natural minerals from the earth manufacturing and chemical engineering technology.

[Edited on 21-9-2011 by symboom]

FrankRizzo - 21-9-2011 at 15:01

Symboom - What Endimion17 was hinting at is that the aluminum is the most important part of this process because it takes **SO MUCH** energy to create the pure metal from the bauxite ore initially. The amount of hydrogen that you'll extract and convert back to usable energy from this aluminum process is a **fraction** of the energy it took to create the metal. You would be using the aluminum metal as a very inefficient battery, charging it up at the mfg plant, and discharging it with your chemical process at a huge loss.

It would be more more efficient to take a semi load of lithium ion batteries to a power plant, charge them, ship them to their final destination, then use that energy to power whatever you need.

not_important - 21-9-2011 at 17:39

Note that the energy loss in going from elemental Al to H2 to some sort of work is fairly high. There's some loss in generating the H2, as low grade heat is also released. Then turning the H2 into useful energy has losses, heat engines are always below the Carnot limit, while low temperature fuel cells do little better once all their overhead is factored in. You'd do better using the Al in a battery, or picking some other metal that could be used in a battery; the efficiency will be significantly higher.

There have been proposals to "export" energy from sunny locations such as the Sahara or Australia via the production of elemental Al, Si, or B, then transporting that to locations where needed.

IrC - 21-9-2011 at 20:42

Quote: Originally posted by symboom

No idea how easy to do in the way you are wanting or the scale. I never get it all back but I also add Indium. I have tried adding other metals but Ga/In works better insofar as getting the Al to react with water goes. For a while I toyed with novel methods to produce hydrogen for a fuel cell in a steady flow by dripping water onto my alloy ingots. A messy waste of time but it sure did give a nice controllable H2 source. This after giving up on electrolysis to get steady hydrogen but in the end I got tired of the hydroxide mess I was cleaning up all the time. Plus recovering the Gallium was a loss ridden pain in the ass. Not something I would consider for a power plant. I went back to the old Zn in acid route, easier to clean up and cheaper. However if you are trying to build a power source all of these methods leave little to be desired. Plus you are never going to get ahead energy wise. If I could I would go with a Pu-Tourmaline generator barring inventing a real Mr Fusion.

symboom - 21-9-2011 at 21:00

Quote: Originally posted by FrankRizzo

Symboom - What Endimion17 was hinting at is that the aluminum is the most important part of this process because it takes **SO MUCH** energy to create the pure metal from the bauxite ore initially. The amount of hydrogen that you'll extract and convert back to usable energy from this aluminum process is a **fraction** of the energy it took to create the metal. You would be using the aluminum metal as a very inefficient battery, charging it up at the mfg plant, and discharging it with your chemical process at a huge loss.

It would be more more efficient to take a semi load of lithium ion batteries to a power plant, charge them, ship them to their final destination, then use that energy to power whatever you need.

well yes i know it takes more energy to make the aluminum like i said.
just experimenting with manufacturing tech. i needed some hydrogen
and i got scrap aluminum and used it and then plan to recycle it back to aluminum with a dc power im still waiting for the lens but that will save me from using so much power just to melt it. my arc welder could do it but that sucker goes through a lot of power

and the lens should work just as good to melt it than use my small industral dc power supply.

symboom - 21-9-2011 at 21:04

Gallium is more reactive than aluminium in many ways. The idea your gallium stays behind unaffected is a complete nonsense.

what were you using Gallium is more reactive than aluminium thats nonsense. aluminum without the oxide disolves in water i am putting it in water

IrC - 22-9-2011 at 03:01

"The idea your gallium stays behind unaffected is a complete nonsense." - symboom

The idea? How about the experience. I make an ingot that looks like a piece of solid cast Al, by thoroughly melting Al, Ga, and In while mixing well in a molten state. After it is a cool solid piece of metal which I can hold in my dry hand, wiping a wet paper towel across it produces steam from the wet paper and the piece of solid metal gets so hot I have to drop it. I put it in water and H2 comes bubbling furiously off, leaving a crumbling hydroxide and/or oxide, producing so much heat steam blows off. The leftover bits are mixed with globs of bright silver looking liquid metal which I can separate out. You call it nonsense I call it experience I have done this with over a pound of my starting material. If nonsense please tell me what the shiny liquid metal is I have left over starting with H2O, Al, Ga, In? To say 'idea' indicates to me that to you it's all theory in your mind you have never actually done the experiment. You are the one coming here asking questions of others please do a thing before you say 'nonsense' to someone who has actually 'done' a thing.

Edit to clarify who I was quoting.

[Edited on 9-22-2011 by IrC]

condennnsa - 22-9-2011 at 03:26

Why not just go the NaOH solution and aluminum way?
Sodium hydroxide can eat through an amazing amount of aluminum, especially if concentrated.
I remember I once tried this and after a while aluminum hydroxide actually precipitates out , and aluminum metal keeps on reacting regardless how much metal I added. I eventually stopped adding because the 'solution' became very muddy.

you could use a gas generator setup with aluminum chips or whatever in the bottom and conc NaOH on top, thus controlling the gas output... i should try this sometime.

blogfast25 - 22-9-2011 at 04:18

Quote: Originally posted by IrC

To say 'idea' indicates to me that to you it's all theory in your mind you have never actually done the experiment.

That's correct. I didn't do this experiment, I've no gallium. What you reported struck me as unusual and I was wrong. Quite amazing that an Al/Ga alloy is so reacive to water, yet leaves behind the Ga, unreacted. But I see you comment on things w/o 'having done the experiment' too, so that's a bit pot and kettle, don't you think?

The notion that synboom will use this to make 'economical hydrogen' by recovering the Al with FCC remains a flight of fantasy though. Read his other threads...

[Edited on 22-9-2011 by blogfast25]

blogfast25 - 22-9-2011 at 04:45

Quote: Originally posted by condennnsa

Why not just go the NaOH solution and aluminum way?
Sodium hydroxide can eat through an amazing amount of aluminum, especially if concentrated.
I remember I once tried this and after a while aluminum hydroxide actually precipitates out , and aluminum metal keeps on reacting regardless how much metal I added. I eventually stopped adding because the 'solution' became very muddy.

you could use a gas generator setup with aluminum chips or whatever in the bottom and conc NaOH on top, thus controlling the gas output... i should try this sometime.

You must have used extremely concentrated NaOH for Al(OH)3 to drop out because NaAl(OH)4 is very soluble. Unless lots of CO2 had gotten to your solution, that converts the aluminate to carbonate, with the oxide dropping out...

IrC - 22-9-2011 at 12:00

OK now I'm confused. I was quoting symboom in my last post not blogfast25.

Melgar - 24-9-2011 at 10:58

For generating hydrogen from aluminum economically, the cheapest way to do it is muriatic acid. It has to be diluted down a lot (like half water, half muriatic) though, or else your "hydrogen" gas ends up containing a whole lot of HCl gas. You can't eliminate the HCl entirely, so if this is a problem, the next cheapest way is with NaOH. H2SO4 doesn't work that well though, since some of it is reduced by aluminum to SO2.

You can use gallium to turn aluminum into aluminum oxide/hydroxide, but if you want to ever see your gallium again, it'd better be pure aluminum, (most structural aluminum is in the form of an alloy) or you'll wind up with a lot of crud to separate it from. Most of that crud will dissolve in dilute HCl faster than gallium will, but you'll still lose some gallium. Aluminum foil and wire are both pure aluminum, and you can use these to make your gallium/aluminum alloy to good effect.

Some notes: it doesn't take much Ga at all to make aluminum react with water. A 1% Ga alloy will do just fine. Of course, the more gallium in the alloy, the faster the aluminum will react. Also, gallium melts above 30 degrees celcius. Once your water reaches gallium's melting point, the reaction will go MUCH faster! Possibly too fast, in fact, so it's a good idea to start with warm water so the reaction doesn't end up going out of control somewhere in the middle. As far as recycling the gallium, a really dilute HCl solution works great for cleaning it off. So does NaOH. Keep in mind though, that both will slowly eat away your gallium too.

[Edited on 9/24/11 by Melgar]

blogfast25 - 25-9-2011 at 05:04

Quote: Originally posted by Melgar

For generating hydrogen from aluminum economically, the cheapest way to do it is muriatic acid. It has to be diluted down a lot (like half water, half muriatic) though, or else your "hydrogen" gas ends up containing a whole lot of HCl gas. You can't eliminate the HCl entirely, so if this is a problem, the next cheapest way is with NaOH. H2SO4 doesn't work that well though, since some of it is reduced by aluminum to SO2.

[Edited on 9/24/11 by Melgar]

Do you have evidence for that? I’ve dissolved Al in fairly strong (up to 50 %) H2SO4 many times and will testify that it is harder to do than with HCl or NaOH/KOH and that it is quite a smelly activity. But I’ve never identified the smell as part SO2. I could be wrong on that, of course…

Melgar - 25-9-2011 at 09:44

I remember reading somewhere that in low concentrations sulfuric acid acts more like a typical acid toward metals, forming the sulfate salt and hydrogen, and that at higher concentrations it acts more like an oxidizer, forming water, metal oxides, and SO2. And also, that you can't really supress either reaction entirely. But sulfuric acid is definitely an oxidizing acid, which is why you get some elemental bromine when you mix it with a bromide salt.

blogfast25 - 25-9-2011 at 11:48

Quote: Originally posted by Melgar

I remember reading somewhere that in low concentrations sulfuric acid acts more like a typical acid toward metals, forming the sulfate salt and hydrogen, and that at higher concentrations it acts more like an oxidizer, forming water, metal oxides, and SO2. And also, that you can't really supress either reaction entirely. But sulfuric acid is definitely an oxidizing acid, which is why you get some elemental bromine when you mix it with a bromide salt.

Yes, no one disputes it is an oxidising acid but with metals that are easily oxidised (like Al) I'd have expected all oxidising to be done by H3O+. I'll be keeping an eye and a nose out next time I dissolve some Al in faitly strong H2SO4...

Melgar - 25-9-2011 at 14:24

Quote: Originally posted by blogfast25

Quote: Originally posted by Melgar

Well, the smell is definitely a sulfur smell, so the H2SO4 is being reduced to something, for sure. There could even be some H2S formed as part of the reaction, if Al is a strong enough reducing agent. Anyway, the reaction isn't a clean one, producing several different products, and is certainly not an ideal way to generate hydrogen. :p

FrankRizzo - 26-9-2011 at 12:57

Quote: Originally posted by IrC

"The idea your gallium stays behind unaffected is a complete nonsense." - symboom

The idea? How about the experience....[snip]

Quote: Originally posted by IrC

OK now I'm confused. I was quoting symboom in my last post not blogfast25.

IrC - No, you were quoting blogfast25. Symboom was simply quoting blogfast25's post, then replying in the line below, basically agreeing with you. He didn't use the quote feature, which made things confusing. Also, his English and punctuation could use a bit of work, but the gist was there.

Quote: Originally posted by blogfast25

Gallium is more reactive than aluminium in many ways. The idea your gallium stays behind unaffected is a complete nonsense.

YAAaawwnn...

[Edited on 20-9-2011 by blogfast25]

Quote: Originally posted by symboom

Gallium is more reactive than aluminium in many ways. The idea your gallium stays behind unaffected is a complete nonsense.

what were you using Gallium is more reactive than aluminium thats nonsense. aluminum without the oxide disolves in water i am putting it in water

AJKOER - 19-11-2018 at 06:03

If a product requires more energy to produce from currently available paths, than it does not make economic sense to manufacture. However, the key phrase here is ‘currently available’ as for limited time spans and setup costs, one may be able to obtain ‘free’ energy. Examples of the latter could include electricity generated from the sun or tides or dams or nuclear power plants or drilling/transporting oil/gas or by tapping in natural occurring heat sources (like volcanoes,...).

As the cost of tapping into 'free energy' based on technological process changes over time, one cannot assert that a product that normally requires more energy to produce from current conventional paths is not deemed forever to be economically unfeasible.

[Edited on 19-11-2018 by AJKOER]

Sulaiman - 19-11-2018 at 06:44

In a global sense the above is true,
but I personally get nothing from my used aluminium that I put in the recycling bin,
so to me (and most SM members) aluminium is a FREE resource.

AJKOER - 19-11-2018 at 07:07

As the value of recycled Aluminum is baked into the sale price of Al, if everyone (or enough of us so that the cost of collecting makes it unecomical) stopped recycling due to newly invented alternative uses, then the current sale price of Aluminum increases for everyone!
---------------------------------

Here is my suggested more fruitful path to hydrogen, the electrolysis of methanol! Apparently, the energy required is but a fraction of the energy required to liberate H2 from H2O (see https://www.techbriefs.com/component/content/article/tb/tech...)!

To quote another source (https://www.researchgate.net/publication/288690690_Hydrogen_... ):

"An economical method to produce hydrogen by direct electrolysis of methanol is developed. It is demonstrated that the hydrogen produced by the electrolysis of methanol can considerably reduce the consumption of electricity. The novelty of this technique is the inherent simplicity and the substantially lowered cost. Using these modified proton exchange membrane fuel cell membrane electrode assembly (PEMFC-MEA) as an electrolyzer, any scale requirements can be easily achieved. The combination of this electrolyzer concept with the concept of solar cells would economically produce hydrogen for storage and subsequent use in, or for in situ use in, fuel cells and chemical engineering applications."

Still not a totally great idea given the energy needed to make the CH3OH.

[Edited on 20-11-2018 by AJKOER]