Sciencemadness Discussion Board

Easy sulfite ion in a pinch...

semiconductive - 23-11-2025 at 18:55

I have sodium sulfite, sodium metabisulfite, and sulfamic acid (no sodium).
I don't have any free source of sulfite ions that don't have a strong alkyl or amino group attached to them.

I got to thinking, maybe I can remove the sodium atoms in a clean way?

eg: Sodium Sulfite salts dissolve only slightly in alcohol and sulfamic acid is soluble in warm alcohol (etoh < 100 C).

But I also read that sulfite ion is very soluble in alcohol by itself.

Could I make free sulfite ions in alcoholic solution by adding finely powdered sodium sulfite or metabisulfite, to a warm beaker of (50 C) denatured alcohol that has 5% sulfamic acid dissolved in it?

I'm thinking sodium sulfamate is insoluble in alcohol, but both sulfamic acid and sulfite are soluble. So, if I add not quite enough sulfamic acid to replace all the sodium in the sulfites; I should end up (eventually) with free sulfite ions in alcohol and a mixed precipitate. (perhaps after chilling to precipitate out as much sodium as possible).

Or would sulfamic acid act as a catalyst, even at 50C temperature, and cause the sulfite ions to react with the alcohol (which sort of defeats the purpose by producing water and an ester).

Is there an easy way to check / test the result ?


[Edited on 24-11-2025 by semiconductive]

DraconicAcid - 23-11-2025 at 20:10

Quote: Originally posted by semiconductive  
I have sodium sulfite, sodium metabisulfite, and sulfamic acid (no sodium).
I don't have any free source of sulfite ions that don't have a strong alkyl or amino group attached to them.


Sodium sulphite does not have an alkyl or amino group attached to it. What are you trying to do?

If you react sulphite with an acid, you get sulphur dioxide, which isn't useful in most circumstances.

semiconductive - 23-11-2025 at 21:22

Sigh. alkaline metal (sodium). It's a typo.
In the text I clearly stated what I'm trying to do.

Sulphur dioxide, when dissolved in alcohol, can be ionized according to AI searches.
If I start with a salt (sulfite), then presumably a non-colloidal looking solution in ethanol has molecules of sodium sulfite or (at least) , or of 2 x Na+ and SO₃²⁻ ions at most, in the case of metabisulfite -- I'm not sure what ion I'll get, but it will have sodium and a sulfur oxide mixture of some kind.

Can I remove the sodium ions while keeping sulfite ions in solution, without it reducing to an ester that isn't ionized ?

AKA: Make a conductive liquid in alcohol, such as I use in my electrodeposition of Nickel thread.
Thanks.



DraconicAcid - 23-11-2025 at 21:51

You cannot have sulphite ions in solution, regardless of the solvent, without a counterion. You could replace the sodium ions with potassium ions, but you can't get rid of the cation completely.

bnull - 23-11-2025 at 21:53

Do not trust AI. It tends to select sentences that conform to positions it has defined beforehand. I don't know how this works but that's what I have seen so far. It also provides sources even when they contradict said position. ChatGPT is very good with translations, and I suppose Grok would do the same if it wasn't for its stiff right arm. Gemini is like a confused elderly relative. They only provide information, not knowledge, and a good chunk of it is wrong.

[***]


Sulfur dioxide is soluble in ethanol. As far as I know, it doesn't ionise there. If AI said it does, ask for the sources and carefully check them out one by one. But beware: I have seen non-existent articles from non-existent publications and authors given as sources by ChatGPT.

Edit: You're not using absolute ethanol, you're using the azeotrope, so perhaps, only perhaps, you may get sulfite ions because of the reaction of sulfur dioxide with water. But sulfurous acid may be strong enough to catalyze esterification to diethyl sulfite.

[Edited on 24-11-2025 by bnull]

semiconductive - 23-11-2025 at 22:07

Yes.. Yes... balance of charges. Very helpful.

To be extra clear:
Alkyl groups and Alkaline metals are both cations.

I don't want the sodium cation.

Which is why I was thinking about getting the sodium to precipitate with the sulfamic acid ion to become insoluble sodium sulfamate ( perhaps in very cold alcohol ).

An ioninzed (lysed) alcohol, ethanol, or hydrogen, as a replacement ion for sodium is fine in my book. If the alcohol (ethanol) is reduced to an alkane without one hydrogen because the oxygen goes away .... I'm totally happy and don't care. I'd say "yaay."

I'm just not happy with ester where *both* of the sulfite charges are neutralized.
I want an *ion*.

I'm in a semi-conductive ion mood right now.

Because I don't have sulfur dioxide gas, I don't think sulfur dioxed is ionized, and I don't know how to make it ionized at 10:00PM my time as I prepare to go to sleep without poisoning myself (again.).

I just want to steal some already ionized sulfite or maybe S₂O₅²⁻ (if that exists as a true radical in solution), AKA: to experiment with sulfer oxide an-ions. (for they are attracted to the anode).

Then I can have sweet dreams.

:cool:

bnull - 23-11-2025 at 22:41

Quote:
I'm just not happy with ester where *both* of the sulfite charges are neutralized.
I want an *ion*.

There's no ethyl bisulfite. It is all or nothing in this case.

semiconductive - 23-11-2025 at 23:47

duplicate post deleted. Internet problems.


[Edited on 24-11-2025 by semiconductive]

semiconductive - 23-11-2025 at 23:49

Quote:
Edit: You're not using absolute ethanol, you're using the azeotrope, so perhaps, only perhaps, you may get sulfite ions because of the reaction of sulfur dioxide with water. But sulfurous acid may be strong enough to catalyze esterification to diethyl sulfite.


The internet is really bad here ( Nov 2025 ) -- the sciencemaddness forums have been inaccessible for hours at a time every other day over a week; but the opening page to the site is just fine. I don't get it.

Note:
I do have access to expensive absolute alcohol, but I'm afraid to open the cap because --- it won't be absolute anymore.

The hardware store stuff is cheap, contaminated with a bit of methanol, and I can put twice baked at 350 [C] magnesium sulfate salt into it which sinks to the bottom. After first baking, I powder it the MgSO₄, after second baking ... I'm convinced it's dry.

I'm pretty sure the resistivity of the ethanol will rise above what my meter can read (20 mega ohm) once a bunch of salt is sitting at the bottom. So, I'm guessing there won't be much water left. I generally do try to get rid of water ....

I will repeat experiments with the expensive stuff, after making my obligatory stupid mistakes multiple times with the cheap stuff. 1 gram of Mg·SO₄ (anhydrous), 3 [CC's] denatured alcohol. I always put kerosene on top to keep moisture out as much as possible...

Yes, Bnull, water is always a problem.

But I think I can at least count on the common ion effect to reduce solubility of sulfates as much as possible ?
maybe not ?

bnull - 24-11-2025 at 01:06

Quote:
The internet is really bad here ( Nov 2025 ) -- the sciencemaddness forums have been inaccessible for hours at a time every other day over a week; but the opening page to the site is just fine. I don't get it.

It's not the internet. It is AI. The leeches scrape the forum a few times a month, leaving the forum barely functional.

Absolute alcohol may prove to be utter useless in this case. You need a little water to make (bi)sulfite, which is in equilibrium with dissolved sulfur dioxide. No water means no (bi)sulfite.

An alternative is to use a cation that is inert during electrolysis. Quaternary ammonium, for example. I don't know if it forms a sulfite or bisulfite.

semiconductive - 24-11-2025 at 01:43

Quote:
There's no ethyl bisulfite. It is all or nothing in this case.


Hmmm.... my thinking:

Sulfite is a triagonal bi-pyramidal ion. It has a resonance structure, which means that the 'oxygens' carrying the two negative charges can shift around the molecule pretty much instantaneously.

During bonding to a cation, one of the two negative charges of the anion will become fixed (neitralized) in a specific oxygen bond; , which means the remaining negative ion charge either remains in resonance among two oxygens or else something (unknown) disrupts the symmetry and one oxygen becomes will more polar negative than the other.


An alcohol can loose either a hydrogen, or the hydroxide group as a whole.
( I've never understood which is more likely, or why. )

Presumably, if the hydrogen leaves then the alcohol is acting as a proton donor (acid). If the hydroxide leaves, the alcohol acts as a base and becomes an alkyl group; R-⁺ + OH⁻ = C₂H₅⁺ + OH⁻

I know Draconic Acid mentioned some years ago that hydrogen doesn't leave by itself; and I'm thinking, without the presence of water, hydronium molecules aren't going to form easily.

But, even then, It is possible to imagine two ethanol molecules to collide and one of them loose an 'H' while the other looses an 'OH', thus giving a temporary situation of R-O-⁻ + H₂O + R⁺.

In that case, the very presence of the water molecule is what prevents the system of two ions from being an ester immediately. ( Shortly thereafter, it might become one if the water is removed or kinetically leaves due to heat. )

If the two alcohol collision arose with a sulfite ion nearby , I don't see why the R⁺ would not be attracted to it while the R-O-⁻ was repelled by it.

I don't quite get why would a single positiviely charged R⁺ ion would not be attracted to a sulfite ion (-2) ?

I'm not asking that an ethyl bisulfite molecule be isolated from the solution. I'm fine if it's a so called 'phantom' molecule. I merely asking why this 'phantom' molecule can't exist as a loose association of ions that is never isolatable.

eg:
A wandering sodium ion, when it does hit a sulfamic acid anion ( negatively charged ) , will form a very stable structure; eg: otherwise, it would be easily soluble in alcohol -- and -- well, it isn't. So sodium sulfamate has to be pretty stable compared to alcohol.

I'm thinking, the big issue here is whether the stability of the sodium sulfamate molcule is enough to remove the sodium from solution or not. Isn't this just a matter of solubility ? The less soluble it is, the stronger the bond must be ?

Please elucidate, what exactly prevents a liquid ethyl bisulfate from existing in solution (dissolved only) if sodium is removed by sulfamic acid precipitation.

Your knowledge is beyond mine, or I've forgotten something.








DraconicAcid - 24-11-2025 at 07:48

Hmmm.... my thinking:
Sulfite is a triagonal bi-pyramidal ion. It has a resonance structure, which means that the 'oxygens' carrying the two negative charges can shift around the molecule pretty much instantaneously.

The charge is delocalized around the three oxygens. It's not shifting.
During bonding to a cation, one of the two negative charges of the anion will become fixed (neitralized) in a specific oxygen bond;

No. The anion does not form a covalent bond to the cation.

An alcohol can loose either a hydrogen, or the hydroxide group as a whole.
( I've never understood which is more likely, or why. )

Presumably, if the hydrogen leaves then the alcohol is acting as a proton donor (acid). If the hydroxide leaves, the alcohol acts as a base and becomes an alkyl group; R-⁺ + OH⁻ = C₂H₅⁺ + OH⁻

Again, no. The alcohol can act as an acid and lose H+, or it can act as a base and gain H+. ROH + ROH <==> RO- + ROH2(+). The equilibrium constant for that reaction is very small (several orders of magnitude lower than the autoionization of water). While one could imagine the two ions then reacting to give an ether and a molecule of water, it's not actually going to happen.

I don't quite get why would a single positiviely charged R⁺ ion would not be attracted to a sulfite ion (-2) ?

They would be, but you're not going to form any extremely electrophilic and unstable carbocations in alcohol solution.

I'm thinking, the big issue here is whether the stability of the sodium sulfamate molcule is enough to remove the sodium from solution or not. Isn't this just a matter of solubility ? The less soluble it is, the stronger the bond must be ?

What cation are you going to replace the sodium with? The hydrogen ion from the acid? In that case, you're going to get H2SO3, which will decompose to give sulphur dioxide.[color]

semiconductive - 24-11-2025 at 11:53

Quote:
The charge is delocalized around the three oxygens. It's not shifting.


Um. I'm not a big fan of the "Wanted: Schrodinger's cat both dead and alive." interpretation. I find it often makes people claim to be impossible some some things which are normally possible.

eg: When a sodium ion Na⁺ approaches a sulfite ion (2-), if the delocalized charge had to *Stay* delocalized, then I suspect the only place the sodium ion could approach the atom would be the top or bottom of the trigonal pyramid ( symmetrically spaced from all the charges ).

Either the charge can shift or it can't, or maybe it's both shifted and not shifted??

For the bonds that eventually happens could include sodium near one of the oxygens -- and not necessarily sodium at the top of the pyramid and equi-distant from all the oxygens.

I'm trying to be inclusive of possibilities rather than exclusive, when I don't know enough to be sure.
Quote:

No. The anion does not form a covalent bond to the cation.


I'm thinking: Covalent vs. Ionic is a matter of degree.
So, you've re-enforced the notion that the anion charge remains de-localized even when a sodium cation is in very close proximity to one of the sulfite ion's oxygens while still being relatively far away from the Sulfur atom.

So, you're giving me new data. (for me).

Is this a 'totally' no change in delocalized charge denisty -- or is it a shifty 1% change which is not covalent -- but still, not zero ?? ( How would I know? )
Quote:
Again, no. The alcohol can act as an acid and lose H+, or it can act as a base and gain H+. ROH + ROH <==> RO- + ROH2(+).


Ok, let's correct my misconception; for this may help me in the future make better guesses: Before, I said I don't know if there is an analog to the hydronium ion, but you seem to be saying there is an alochol analog to the hydronium ion. It's ROH₂⁺. Correct ?
Quote:
The equilibrium constant for that reaction is very small (several orders of magnitude lower than the autoionization of water). While one could imagine the two ions then reacting to give an ether and a molecule of water, it's not actually going to happen.


Again, I wasn't even trying to say an isolatable ether gets formed.
But, I think you're giving me quantitative argument and not a qualitative one.

I don't have quantiative data -- especially when not talking grossly aqueous soltuions ; and I don't have your experience.

eg: As far as I know -- The relative re-ionizable magnitude of sodium sulfamate precipitate in alcohol may also be orders of magnitude smaller than autoionization of water. But, if the ability to ionize the precipitate exists at all and is smaller than that of the ether reaction -- the reaction could proceed ( but possibly very slowly).

This is the kind of thought that was going through my head before falling asleep last night.

AKA: Without your data (with no general citation) -- I have no rule of thumb to estimate the relative liklihood or make better predictions in the future. Which I would like to be able to do.
Quote:

What cation are you going to replace the sodium with? The hydrogen ion from the acid? In that case, you're going to get H2SO3, which will decompose to give sulphur dioxide.


I'm not sure of your point.
I'm experimenting, I will try many things.

I know H₂SO₃ does not exist in isolation; but I was not trying to isolate it.
H₂SO₃ has a neutral leaving group of H₂O -- and it follows that the molecule could split into SO₂ + H₂O.
I Agree.


But -- Is the mere presence of hydronium ions in solution enough to cause a nearby sulfite ion to decompose and leave solution ?

For, then SO₂ gas ought to be produced in proportion to the probability of the presence of hydronum squared in all sulfite solutions. I don't smell much sulfur dioxide and I am not sure how big of a number I should assign to that as an estimate.

I'd like to learn how to predict based on what I can measure or detect at home.















semiconductive - 24-11-2025 at 14:04

Bnull:

One of the top three chemicals I thought of try as an 'cation' replacement was Choline, isolated from choline chloride. I think that's a quaternary ammonium compound. I have a bag of it. So, yes... that's possible.

It's the same bag I bought to try choline chloride + Urea that Draconic recommended -- and which turned black instead of plating.

I've got all of urea, thiourea, choline chloride, oxalic acid, and a computer controlled thermometer with soldering iron and something called a 'schiff base' sitting on my desk. I even have a glass coated electrode that can be charged to +1126volts DC in order to attract anions in a thin layer near it's surface.

Sorry, I've been trying to post this partial answer for over an hour. The AI stuff is basically starting to lock me out ... I can't even read the site other than the login page. I'm quite frustrated right now! :(

chornedsnorkack - 24-11-2025 at 15:26

Quote: Originally posted by bnull  


Absolute alcohol may prove to be utter useless in this case. You need a little water to make (bi)sulfite, which is in equilibrium with dissolved sulfur dioxide. No water means no (bi)sulfite.

"Bi"sulfite or hydrogen sulphite?
Hydrogen sulphite has extra decay paths - more options to eliminate compared to sulphite
C2H5OH+SO2 <-> C2H5OSO2H
but I suspect the equilibrium would be on the left, towards elimination of SO2
Quote: Originally posted by bnull  


An alternative is to use a cation that is inert during electrolysis. Quaternary ammonium, for example. I don't know if it forms a sulfite or bisulfite.

With excess of SO2, I suspect "bi"sulphite. But the equilibrium of these would depend on solubilities.

bnull - 24-11-2025 at 19:41

Quote:
"Bi"sulfite or hydrogen sulphite?

Both are the same. And "(bi)sulfite" is shorthand for "bisulfite or sulfite or a mixture of both".

I don't know if $$C_2H_5OH+SO_2 \leftrightarrow C_2H_5OSO_2H$$ happens (zero indication so far and the only possible source I found is paywalled). I was thinking of $$SO_2+H_2O \leftrightarrow H_2SO_3.$$ Sulfurous acid is unstable and decomposes to sulfur dioxide and water. Alcohol would work as a dehydrating agent in this case.

What I don't know is if quaternary ammonium forms bisulfite or sulfite. Maybe both, maybe one of them, maybe none depending on the specific radicals. I'm in the dark here.

[Edited on 25-11-2025 by bnull]

chornedsnorkack - 25-11-2025 at 00:24

Quote: Originally posted by semiconductive  

To be extra clear:
Alkyl groups and Alkaline metals are both cations.

Alkyl cations are very hard to get and very active.
Quote: Originally posted by semiconductive  

Which is why I was thinking about getting the sodium to precipitate with the sulfamic acid ion to become insoluble sodium sulfamate ( perhaps in very cold alcohol ).

You´re considering very cold alcohol?
The liquid range of neat ethanol at 1 bar is from +78 to -114.
The liquid range of neat sulphur dioxide also at 1 bar is from -10 to -75.
Quote: Originally posted by semiconductive  

Because I don't have sulfur dioxide gas, I don't think sulfur dioxed is ionized, and I don't know how to make it ionized at 10:00PM my time as I prepare to go to sleep without poisoning myself (again.).

I just want to steal some already ionized sulfite or maybe S₂O₅²⁻ (if that exists as a true radical in solution), AKA: to experiment with sulfer oxide an-ions. (for they are attracted to the anode).

Then I can have sweet dreams.

:cool:

You need a suitably polar solvent.
At about 20 Celsius, the dielectric permittivity of ethanol is 25.
At -10, the dielectric permittivity of sulphur dioxide is 16.
Not awfully good for ions but not quite intolerable either.
Certainly sodium ethoxide C2H5ONa has high solubility in ethanol (20%). Do sodium ethoxide solutions in dry ethanol conduct electricity and electrolyze?
Would sodium ethoxide react with dry acidic oxides? Like
C2H5O-+SO2=C2H5SO3-
C2H5O-+CO2=C2H5CO3-?

bnull - 25-11-2025 at 02:51

Liquid sulfur dioxide dissolves quite a few salts, especially quaternary ammonium salts. See p. 40 of Waddington, Non-aqueous Solvent Systems (can be borrowed at https://archive.org/details/nonaqueoussolven0000unse).
Quote:
Would sodium ethoxide react with dry acidic oxides? Like
C2H5O-+SO2=C2H5SO3-
C2H5O-+CO2=C2H5CO3-?

The second reaction is well known, it is a way to make alkyl carbonates. The first reaction is the problem. As far as I could find, reactions involving sulfur dioxide and alkoxides result in dialkyl sulfites, which do not produce ions. If it happens the way you wrote, then sulfur dioxide is essentially lost as the alkyl sulfonates are stable. Esylic acid (ethanesulfonic acid) is strong and stable.
Quote:
Certainly sodium ethoxide C2H5ONa has high solubility in ethanol (20%). Do sodium ethoxide solutions in dry ethanol conduct electricity and electrolyze?

S. Tijmstra wrote a paper about the conductivity of sodium methoxide and ethoxide in alcoholic solutions for the Zeitschrift für physikalische Chemie, Volume 49, beginning from page 345 (https://archive.org/details/sim_zeitschrift_physikalische_ch...). My German is not that good and I'd probably miss a few words and mistranslate the whole thing.

Edit: Typo.

[Edited on 25-11-2025 by bnull]

DraconicAcid - 25-11-2025 at 11:53

You might find sodium lauryl sulphonate (a common detergent ingredient) to be sufficiently soluble in alcohols.

semiconductive - 25-11-2025 at 13:23

Quote:
The liquid range of neat ethanol at 1 bar is from +78 to -114.
The liquid range of neat sulphur dioxide also at 1 bar is from -10 to -75.


I can get dry ice at the local store. That goes to -78.5 [°C].
I've tried it on ethanol and acetone before and got them in the -40's, I seem to recall them becoming slightly hetrogenous. But, that might be because I didn't dry them before cooling.

Hmmm..

Water becomes ice which is a solid. But, it's highest density is at 4 [°C].
The spacing of atoms has to move very slowly in both situations. But, that means the math will be very close (qualitatively) to a solid making the QM/Boltzmann math identical to semiconductors physics.

I'm going to try computing the auto-ionization of water from ice data, and see if I get an accurate approximation. If I do, then I'll assume alcohol is less complicated (because of larger masses), and repeat the same procedure to approximate the properties of alcohol.

----

In solid state semiconductors the intrinsic carrier concentration function is well known:

Ni = C₁ · T^(3/2) · e^(- E/( 2·k·T ))

The energy gap (E) is a smooth function that decreases with temperature and usually has a linear and an Arrehnius factor in it. ( A rational polynomial can easily model it given three data points. )

E ≈ a·T / ( c + b/T )

Looking up a bunch of data points from different authors, I curve fit a
linear log model of water auto-ionization near the freezing point of water:
(I have no way of knowing how accurate this is).

Bounding Kw ⪝ 3.008·( 3.661 - 1000/T )-14.94

Converting the intrinisic carrier concentration to a log10 formula (like pH):
Ni = C₁ · T^(3/2) · e^(- E/( 2·k·T ))
Assume E is in electron volts, just like a semiconductor:

Ni = ( Const + T^(3/2·ln(T) - E/( 2·k·T ) ) / ln(10)
log₁₀( Ni ) ≈ ArbitraryConst + 0.6514*ln( T ) - E·2553/T

The energy gap between H⁺ and OH⁻ should just be the difference in ionization potentials between neutral water and released ions.

Since hydrogen gas electrode is the standard reference at 0V for oxidation/reduction tables; I think I can get away with the E = energy required to convert H₂ gas with aqueous hydroxide ions into water.

so Eg ≈ 0.83 volts at 25 [°C].
If I'm wrong we'll soon know:

2553·0.83 ≈ 2119

A crude theoretical calculation just assumes Energy gap (E) is constant from freezing to standard temperature 25 [°C]:

At freezing:
Kw ≈ -14.94
Kw = log₁₀( Ni ) ≈ ArbitraryConst + 0.6514*ln( 273.15 ) - 2119/273.15
-14.95 ≈ ArbitraryConst + 3.6543 - 7.7576
ArbitraryConst ≈ -10.85

Therefore, I have created a crude model for temperaturem in celcius:
crude Kw ≈ -10.84 + 0.6514·ln( 273.15 + Tc ) - 2119/( 273.15 + Tc )

This is slightly lower than the upper bound calculation. I am confident that the math is qualitatively correct. The slope is lower but not even off by half, so I'm confident I can compute a refined model that will fit very well.

A refined model will modify the values of the first term (constant) and the last term (Arrhenius physics) to get the correct slope of ionization ; and my experience with solids is that extrapolation is usually pretty accurate at higher temperatures. If it's also accurate in liquds we can estimate.

I'm going to just use freezing and standard lab temperature since the standard ionization potentials are known precisely. I'll post a refined equation tomorrow. (see post below.)

DraconicAcid - 25-11-2025 at 15:22

For methanol, pK(autoionization) = 22.67
https://pubs.acs.org/doi/10.1021/acs.jpca.5c03979

clearly_not_atara - 25-11-2025 at 15:24

Quote: Originally posted by bnull  
Liquid sulfur dioxide dissolves quite a few salts, especially quaternary ammonium salts. See p. 40 of Waddington, Non-aqueous Solvent Systems (can be borrowed at https://archive.org/details/nonaqueoussolven0000unse).

Quote:
Would sodium ethoxide react with dry acidic oxides? Like
C2H5O-+SO2=C2H5SO3-
C2H5O-+CO2=C2H5CO3-?

The second reaction is well known, it is a way to make alkyl carbonates. The first reaction is the problem. As far as I could find, reactions involving sulfur dioxide and alkoxides result in dialkyl sulfites, which do not produce ions. If it happens the way you wrote, then sulfur dioxide is essentially lost as the alkyl sulfonates are stable. Esylic acid (ethanesulfonic acid) is strong and stable.


I think you are confusing the alkyl sulfonates with the alkyl sulfites. What he hopes to produce is "monoethyl sulfite (-1)". The corresponding hydrogen ethyl sulfite is probably strongly disfavored (sulfurous acid is basically not observed in solution, while carbonic acid is present to a small extent). The SMILES CCOS(=O)O- is a plausible result of SO2 + EtO-, but the following rxn may destroy it:

2 CCOS(=O)O- >> CCOS(=O)OCC + SO3(2-)

But there are just no sulfite ions without counterions, and any variant of this would risk exposure to large quantities of SO2 gas. I am concerned about recommending any procedure to someone who does not understand why.

davidfetter - 25-11-2025 at 16:29

Quote: Originally posted by semiconductive  
according to AI searches.


STOP RIGHT THERE

If you're consulting AI for literally anything, you do not have the judgment needed to mess with chemistry. Doing so is a sign that you need to do some pretty large reassessments of what you're doing with your life, what sources of information you trust, and what you use to establish that trust. Chemistry can be extremely unforgiving, and AI will happily tell you to do things in that field that will kill you and could kill people near you.

bnull - 25-11-2025 at 19:03

Quote: Originally posted by clearly_not_atara  
I think you are confusing the alkyl sulfonates with the alkyl sulfites. What he hopes to produce is "monoethyl sulfite (-1)".

No, I'm not confusing them. What @chornedsnorkack wrote was a reaction between an ethoxy group and sulfur dioxide with ethanesulfonate as product (ethoxy loses its oxygen to sulfur and a bond is formed between sulfur and the carbon). As far as I know, it doesn't work that way. What @semiconductive wants is monoethyl sulfite (ethyl bisulfite), or some source of sulfite ions in nonaqueous medium that do not contain or generates alkaline cations. As far as I know, ethyl bisulfite does not exist. If it did exist, it wouldn't dissociate to ethyl and sulfite ions.

H. F. van Woerden wrote a review about organic sulfites (https://doi.org/10.1021/cr60226a001). I haven't read it yet. Maybe there's something there about monoalkyl sulfites.

As this whole thing has to do with nickel plating in non-aqueous solutions, the question that should be asked is, is nickel sulfite soluble in any non-aqueous solvent? The SDS below has a few references that may lead somewhere.

Attachment: SDS-26-pages_258.pdf (58kB)
This file has been downloaded 53 times

semiconductive - 25-11-2025 at 22:43

@DraconicAcid, Thank you for the Methanol link. That looks very promising as a second test and to build intuition about alcohol vs. water. I wouldn't have been able to find it myself.

General question:

I am looking at a NIST page on standard potentials in water:
https://www.nist.gov/system/files/documents/2019/04/02/jpcrd...

There are a few problems with NIST data, such as no mention of isotope blends, etc., so I can't get really be sure what conditions the experiment was done under. But, this is where I normally get engineering data....

In the NIST paper, I see the hydrogen gas vs. hydroxide ions reaction listed as:
E°=-0.828 , ΔE=-0.0008360 For H₂(g).OH⁻ ⟷ H₂O(liq)

I know from electronics that the energy 'band-gap' in semiconductors is affected by whether the chip is packaged in epoxy and under compression -- or the die is bare and exposed to air ; the same should happen in electrolytes. So, I suppose it's possible to reverse the trends of the energy gap by putting it in pressure container. But, normally I would expect the magnitude of E to get smaller with increasing temperature.

But, the sign convention of the energy change is negative, and the document equation (2) on page (2) shows the coefficient as added to the initial E in proportion to temperature.

That doesn't make sense to me. The analogy seems broken.

Does hydroxide to water conversion create more voltage as the solution gets hotter?

If I naively apply the NIST equation (2) -- I get energy gaps that indicate a voltage magnitude increase:

['-0.8071', '-0.8155', '-0.8196', '-0.8238', '-0.8280', '-0.8322', '-0.8489', '-0.8907']

Therefore:
I grabbed some auto-ionization of water values vs. temperature from online searches, and converted them to exponents and averaged them to compare to do a sanity check.

Tc =[ 0, 10, 15, 20, 25, 30, 50, 100 ]
log₁₀(Kw)=[ -14.94, -14.54, -14.35, -14.17, -14.00, -13.83, -13.26, -12.30 ]

If I use the NIST values with increasing voltage magnitude, I get pretty bad agreement:

Eg = ['-0.8071', '-0.8155', '-0.8196', '-0.8238', '-0.8280', '-0.8322', '-0.8489', '-0.8907'
log Kw= ['-14.51', '-14.30', '-14.19', '-14.10', '-14.00', '-13.91', '-13.56', '-12.86']

On the other hand, if I use the wrong sign for equation (2), the energy gap goes:

['-0.8489', '-0.8405', '-0.8364', '-0.8322', '-0.8280', '-0.8238', '-0.8071', '-0.7653']
['-14.90', '-14.52', '-14.34', '-14.17', '-14.00', '-13.84', '-13.23', '-12.00']

It's clear for small deviations (5 degrees) that using the 'wrong' sign agrees very closely with experimental data that can be found online.

Before I curve fit the correction -- does anyone know why the sign of the potential change is negative instead of positive in the NIST paper? Is it a convention, or a typo?

semiconductive - 28-11-2025 at 12:07

Quote:
As this whole thing has to do with nickel plating in non-aqueous solutions, the question that should be asked is, is nickel sulfite soluble in any non-aqueous solvent? The SDS below has a few references that may lead somewhere.


Actually, I put this in a different thread because it has more general application than just nickel plating. Regarding plating: I am actually more interested in plating iron, and iron pyrite (Fools Gold), than nickel. But: I've been attempting nickel because it's easier to reduce from solution than iron.

Note: I successfully plated grey iron this last week in a solution where iron oxalate ought to have been insoluble. But the test tube is super bright yellow and very conductive. I also succeeded from an acetone bath, and also using di-cyanamide as a complexing agent. So -- I've actually had amazing progress this last month after failures for two solid years.

If you do research on sulfites, I think you will find it is generally going to be the case that sulfites except of sodium, potassium, and ammonium, tend to be hard to dissolve.

However, what I'm finding is that the common solubility rules of (rarely, but with notable exceptions), do not apply when double salts are made.

Aluminum, for example, has oxidation state +3, and therefore can not be totally bonded with just a single sulfite molecule. If I half-neutralized sulfite (or metabisulfite, which we've sort of ignored) using aluminum; then there will be one bond left over which could be occupied by nickel, an alkyl, iron, or other cation.

You're probably very familiar with potassium alum, which is a common chemical to find in nature. It's extremely fond of absorbing water. But, I've done a few experiments in methanol and it will happily absorb methanol in place of water yielding a new gelationous substance that is quite conductive of electricity.

There ought to be similar chemicals that can be made with sulfite or meta-bisulfite, which both have the same -2 maximum charge as ions as sulfate has. ( But, I expect the properties are going to be slightly different -- I have no idea if they will be better or worse candiates, and am just experimenting! )

Since potassium aluminum sulfate in methanol has some plating activity, I wonder about analogs like lithium aluminum sulfite or potassium aluminum meta-bisulfite.

However, the form I have these acids in always have sodium attached to them. Sodium sulfite + hydrochloric acid, is not a good choice!!! So I'm looking for ways to remove the sodium without producing SO₂ gas ...

For the most part, I'm looking for metastasis reactions that allow me to get rid of sodium and replace it arbitrarily.

But, my chemistry knowledge is very limited.


[Edited on 28-11-2025 by semiconductive]

DraconicAcid - 28-11-2025 at 13:54

I wouldn't expect aluminum sulphites to be stable in the presence of water or hydrogen ions.

Al(3+) + 3 HSO3(-) ---> Al(OH)3 + 3 SO2(g)

semiconductive - 28-11-2025 at 16:46

Quote: Originally posted by DraconicAcid  
I wouldn't expect aluminum sulphites to be stable in the presence of water or hydrogen ions.

Al(3+) + 3 HSO3(-) ---> Al(OH)3 + 3 SO2(g)


Indeed.

And there is some odor using sodium aluminum sulfite in water.
It's not a lot of gas, but it's obviously possible for some to escape.

I find it rather curious that sulfite salts are stable at all.

I'm thinking sodium sulfite Na₂SO₃ could be thought of as Na₂O + SO₂, and in the presence of hydronium or 'alk'onium ions, the Na₂O could become hydroxide radicals. As far as 'leaving' groups go, Na₂O is neutral just as H₂O is neutral.

So, I don't really understand why even stable sodium sulfite doesn't absorb water and release SO₂ gas *all the time*.

The meta-bisulfite is less puzzling to me because the oxygens are not easily grouped into Na₂O. I think it's probably a much bigger molecule that would have to 'leave', and that might help keep the SO₂ groups mechanically 'stuck'.

But, I still don't totally get why it doesn't just decompose down in to sulfite and then into SO₂. I have weird dreams that don't actually happen when tried in test tubes.

bnull - 28-11-2025 at 18:14

Quote:
or metabisulfite, which we've sort of ignored

By design rather than by accident. It hydrates to bisulfite in contact with water and I suppose you're going to make your other sulfites using aqueous solutions. You may try using another solvent, although I'm not sure if the metabisulfite ion will remain as it is or will decompose to sulfite plus sulfur dioxide. I'd bet on the latter as I never had heard of, say, aluminum metabisulfite or iron metabisulfite.

I know of only two ways of removing sodium ions. One uses uranyl acetate (plus some zinc or magnesium ions to make the triple (?) salt) and the other uses ion-exchange resins. The advantages of the resins are that they are cheaper than the uranyl salt and reusable.

Edit: One more thing. Did you use a solution of aluminum sulfate?

[Edited on 29-11-2025 by bnull]

semiconductive - 28-11-2025 at 23:39

The aluminum sulfate comes as an anhydrous powder. It can be mixed with ethanol, methanol, or other alcohols to avoid water. Lithium, potassium, or other salts of very weak acids can be added.
Most of my successes have come from using alkalai citrates.
For example, lithium citrate made in water, can be dried in an oven at around 215 farenheight without decomposing. (101 to 102 Celsius spread in a thin layer). Powdered after drying, citrates can be added to aluminum sulfate in an alcohol solution. I've tried many variations of other organic salts. A fair number of them will loose amounts alkalai to the aluminum sulfate under heated conditions. At that point, the solution will start to form a gel and begins absorbing alcohol molecules because (I assume) there is insufficient water for the alum to become hydrated. The alcohol molecule is the next closest thing to water that is available...


I've tried lithium carbonate in methanol with aluminum sulfate powder, but it's nowhere near as effective.

semiconductive - 29-11-2025 at 00:30

--- Continuing on to the auto-ionization of water ---

I don't hear anyone explaining why NIST's sign is different from what I expected; so I'm going to make a guess and move forward.

Looking around, I see in a Wikipedia, a relationship that is empirical and slightly easier to work with than the linear to Arrhenius relationship that can be derived from lump modeling of atoms.

The empirical formula is called "Varshni's" correction to band gap narrowing.

https://en.wikipedia.org/wiki/Band_gap

I think It has sufficient degrees of freedom to handle the phase change effects of ice into water, or even water into pressurized steam.

What I am going to do next is 'wrong'.

I am going to ignore the sign of the formula from NIST, and use my experience from semiconductor design to make chemical predictions.

Disclaimer: Do not use this technique in any professional settings or where safety is paramount. The professional documentation from NIST disagrees with me. I don't know why.

But:

The differences between water and solids is mostly confined to the ability of ion donors to migrate during ionization.

In liquids both hydrogen atoms and electrons can hop from one group of atoms to neighboring ones, and simultaneously the atoms themselves can re-orient or mix.

This extra motion means that fluids have one more degree of motion freedom than solid semiconductors do. The quantum band shape can change with *time* as fluids re-arrange themselves.

Re-arrangement of the 'band' structure also implies that localized regions of pH must change with time even when the chemicals, themselves, are at 'equilibrium' sealed and isolated in a container.
This is something I was never taught in undergraduate classes, but is a necessary consequence of band theory as I understand it from electrical engineering.

The closest analogy I can think of is chemical oscillation, where colored solutions go back and forth between two states several times before settling down to an equilibrium condition. Although the bulk oscillation *appears* to stop, I want to suggest that it still continues at a microscopic level with random changes in color that cancel out on average. The pH shifts with time in a liquid.

Even with pure distilled water, rippling ionization effects must be occurring that semiconductor equations don't model.

With that in mind, I'm going to use the NIST data (ignore the time dependency) and change the sign of the ionization rate to agree with what it would be in solid semiconductors.

I'm merely going to figure out what constants alpha and beta applied to Varshni's correction will yield the same derivative (change in ion concentration vs. temperature) at standard conditions as is reported in the NIST document. But I am going to ignore the SIGN of the slope, which is a definite error on my part.

Then I'll make a chart of auto-ionization strength for distilled water based on the semiconductor analogy.




semiconductive - 29-11-2025 at 12:41

Working out an energy gap correction for distilled water, or very dilute ionic strength solutions.

Any two distilled water solutions from different labs will likely have slight variations in properties. ( Who knows how much deuterium is in any given source of water... )

But, here goes: Data magnitude is taken from NIST, author "Steven G. Bratcsch" -- and I preserve the original sign of ΔE/ΔTc to check the 'correct' calculations before doing my own 'wrong' calculations.

For H₂(g).OH⁻ ⟷ H₂O(liq)
Tc=25 [°C], E°=-0.828 [eV], ΔE/ΔTc=-0.0008360 [eV/°C]

I'm implicitly converting to energy (electron Volts) which is pressure in volts multiplied by electrons involved. Electronic multi-meters generally measure only a pressure in volts.

The intrinsic ion/carrier concentration (Ni) equation from semiconductor physics has the following logarithmic form ( assuming a constant energy gap ). There are two arbitrary constants, A and B, which depend on material properties.

log₁₀( Ni ) = A + 0.6514·ln(T) - |E₀|·B/T

For Isobaric conditions, STP, H₂O:
log₁₀( Ni ) = -13.996 At 25 [°C]

Therefore:
-0.828 [eV] ·B/298.15 [K] + 0.6514·ln(T) = -13.996... - A

At freezing, I find the following data:
chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/16%3A_AcidBase_Equilibria/16.03%3A_The_Autoion ization_of_Water

But, the author does not say if this is a theoretical value or an experimentally measured value of ice-water. ( This is my life... ugh! )

log₁₀( Ni ) ≈ log₁₀( 1.15/10¹⁵ ) = -14.9393...

Considering the trouble I run into when asking for the pH measurement of distilled water even at 4 [°C] from Google, I am going to do a little more research.

I can't find any pH probe measurements at 4C. People don't report them.

If I search for carbon dioxide error in pH measurements at 4[°C], I find notes that the pH is often between 5.5 and 6.5 due to CO₂ gas absorbtion. Taking the more basic measurement as the least CO₂ affected; then log₁₀( Kw ) is between -13.0 and -14.939.

Sigh: The possible range of data represent completely different qualitative changes from standard temperature and -log₁₀( Kw ) = 13.996. It's not just AI that aren't trustworthy, the original documents don't publish useful data that conclusions can be made from. !

I'll try to use conductivity experimental data to isolate another corroborating auto-ionization value at a nearest to maximum density temperature of 3.98 [°C].

https://www.researchgate.net/publication/237310270_The_Funda...

But, the footnote shows he got the Kw data from someone else...
Reference 4: 4. E. Schmidt, Properties of Water and Steam in SI-Units, Springer-Verlag, New York -- 1969

at 0 [°C] -log₁₀( Kw )≈14.9412
at 5 [°C] -log₁₀( Kw )≈14.7287

Considering the fact that I have two values from different authors that are very slightly different; I'm now going to (temporarily) assume both these values are experimentally valid; ( but in reality, I'm in the same dilemma. I do not know how the values were arrived at. ) A voltage probe measurement is physical, and is in volts. An ionization constant Kw is inferred, indirectly or calculated theoretically.

Half cell voltages change close to 'linearly' according to NIST publications with respect to temperature. ( But this comment is rather suspect!!! )

Therefore, I'm first going to solve three equations in three unknowns presuming the energy gap does not significantly change over the small temperature range of 0 to 5 [°C], but does over 0 to 25 [°C]. I want to know what the E gap needs to be (approximately) near freezing to test for linearity.

I do know by voltage measurement what the energy gap is at 25 [°C].

-0.828 [eV] ·B/(273.15 + 25 )[K] + 0.6514·ln(298.15 [K]) = -13.996... - A
-E [eV] ·B/(273.15 + 5) [K] + 0.6514·ln(278.15 [K]) = -14.7287... - A
-E [eV] ·B/(273.15 + 0) [K] + 0.6514·ln(273.15 [K]) = -14.9412... - A

Therefore:
A ≈ -4.99237...
B ≈ 3822.47...
E (near 2.5 [°C] ) ≈ 0.824345...

What this shows is that for all three data points to be experimentally valid, the energy gap near freezing will be smaller than standard conditions by around 4 millivolts.

But: That means the experimental data is nowhere near linear.

The NIST data claims:
ΔE/ΔTc=-0.0008360 [V/°C],
Over a 25 [°C] change, this gives:
E = -0.828 + ( 273.15 - 298.15)·(ΔE/ΔTc ) = -0.8071 [ V ] .

An approximately linear change is around +20 [mV] according to NIST.
That's ~500% larger than the available data supports. ( Bad words omitted. )

---- Thoughts ----

Since the voltage (gap) is with respect to a hydrogen gas electrode, the half cell reaction is also identical to a full cell reaction. I can't mess that up.

The Nernst equation cancels out since auto-ionization is an 'equilibrium' condition by definition.

E°cell = R·T/n · ln( K )
E°cell = 0.0592/298.15 · T · ln( K )/n
E°cell = 0.0001986 · T · ln( K )/n

The NIST publication didn't list a balanced equation for the hydroxide/hydrogen reaction.

But, I think it is: 2·[OH⁻] + H₂(g) ⟺ 2·H₂O(l) + 2e⁻
E°cell = 0.0001986 · T · ln( K )/2

Still, this doesn't advance me. K is obviously temperature dependent. That's an Enthalpy Entropy relationship, which I haven't done in over 30 years ...

The reduction of 2·H₃O⁺ (aq) + 2e⁻ ⟺ H₂ + 2·H₂O is by definition, zero volts.
But -- Aha! -- there is energy stored when ions appear in water with a dielectric constant separated from other ions of opposite charge. There's something subtle going on...

I give up for today.


[Edited on 30-11-2025 by semiconductive]

semiconductive - 30-11-2025 at 10:21

Quote: Originally posted by chornedsnorkack  

Alkyl cations are very hard to get and very active..


Thank you. I'm beginning to understand that.
Quote:

You´re considering very cold alcohol?
The liquid range of neat ethanol at 1 bar is from +78 to -114.
The liquid range of neat sulphur dioxide also at 1 bar is from -10 to -75.


I tried looking up 'neat' alcohol and sulfur dioxide, and got strange definitions saying it ought to be drunk at room temperature.
Quote:

You need a suitably polar solvent.
At about 20 Celsius, the dielectric permittivity of ethanol is 25.
At -10, the dielectric permittivity of sulphur dioxide is 16.
Not awfully good for ions but not quite intolerable either.
Certainly sodium ethoxide C2H5ONa has high solubility in ethanol (20%). Do sodium ethoxide solutions in dry ethanol conduct electricity and electrolyze?
Would sodium ethoxide react with dry acidic oxides? Like
C2H5O-+SO2=C2H5SO3-
C2H5O-+CO2=C2H5CO3-?


I can measure the permittivity of my solutions with a capacitance meter.
I will have to buy a little reagent grade sodium ethoxide and do a test. ( It'll be a week or two... )

The auto-ionization constant of water (Kw) that I'm trying to understand and model in dry ethanol is sometimes estimated using permittivity calculations. I'm not familiar with this technique.

I see key words like: Born equation, and some debate over whether it is Enthalphy or Gibbs free energy related. So, the accuracy of the calculations isn't something i understand yet.

I can also find articles like the following (Which I am slowly reading and absorbing):
https://srd.nist.gov/JPCRD/jpcrd696.pdf

But, perhaps you already know

The relative permittivity of water is approximately:
ε_r ≈ 1.94315 - 0.0019720·( T [K] -273.15 )

I can measure this for neat alcohol and create an equation, easily.
I just need to 3D print a capacitive cell to hold alchohol while freezing it.

I'm thinking:
Hydrogen atoms in vacuum take 13.6 [eV] of energy to completely ionize.

The energy stored in a capacitor is E = 1/2·c·V²

How is the energy required to completely ionize hydrogen in water related to the energy required to ionize electrons in vacuum?

Is there a simple physical relationship that I might exploit?


[Edited on 30-11-2025 by semiconductive]

semiconductive - 30-11-2025 at 10:30

Quote: Originally posted by davidfetter  
Quote: Originally posted by semiconductive  
according to AI searches.


STOP RIGHT THERE

If you're consulting AI for literally anything, you do not have the judgment needed to mess with chemistry. Doing so is a sign that you need to do some pretty large reassessments of what you're doing with your life, what sources of information you trust, and what you use to establish that trust. Chemistry can be extremely unforgiving, and AI will happily tell you to do things in that field that will kill you and could kill people near you.


Yes, I plan on scheduling a third mid-life crisis for next Wednesday at 2. I take anti-anxiety meds and wonder each day when I wake up God wants me to live. I keep telling my counsellor that I'm not exactly suicidal, but if I were to die -- that'd be OK. There's not much to live for when everything I want to do is out of reach.

I appreciate your thoughts:
I'm not worried about myself, but killing my Mom would be a problem.
I am trying to be careful.

semiconductive - 12-12-2025 at 10:54

I'm running a control experiment, today.

I'm using sulfate in place of sulfite, because sulfate is very stable.
I want to see if a a lithium ferric double salt can be made to dissolve in alcohol.

2CC's denatured ethyl alcohol.
4CC's ethyl citrate (esterified) -- reasonably pure 99.9%.
2CC's kerosene as a cap to keep air and moisture out.

Oven dried lithium citrate at 215 [°F] for 4 hours.

Oven dried Ferrous Sulfate Monohydrate. 280F for 4 hours, 475 for 20 minutes.
I used analytical grade heptahydrate. I pre ground it before baking with a glass rod to make a fine dust. It sticks to glass severely after baking, so the grinding has to happen before in order to make a fine powder.

Even though this is baked in air, there is only a very slight darkening of the dust at the lower temperature. It became a very light tan, almost white.

Note:
Stirring and higher temperatures are a mistake, for it noticeably darkens the salt. Probably Oxygen from air reacting with iron to make it Fe(III) rather than Fe(II).

289 [mg] FeSO₄·H₂O
118 [mg] Li₃[ citrate ]

If I've done the math right, I ought to have around 20 molecules of ethanol for every molecule of Ferrous Sulfate. There ought to be one lithium atom for every sulfate atom.

Sulfate ions have a very weak bond for the first ionization, making it nearly completely ionize in water. The second bond is much stronger making the second ionization a weaker acid. It's harder to break the stronger bond, so I'm hoping to half neutralize a fair portion of sulfate anions with lithium. I'm hoping the lithium will occupy the stronger bond, which will leave the weaker bond to hold onto iron. This might make it electroplate better.


Salt and liquid, stirred, makes a colloidal suspension that falls out in a matter of an hour to the bottom of the test tube.

I heated the solution to 80 [°C] for 4 hours and conductivity very slowly rises. (used 12 steel washers as an anode on a nylon insulator).

Less than 1/5th the salt dissolves into solution.

If I raise the temperature to 102 [°C], there is notice-able bumping, but only small bubbles of gas escape. Most of the salt enters solution. Less than 1/2 CC of solution evaporates in 20 hours of heating. The top of the test-tube never rises above 40[°C].

The colloidal suspension returns and remains for as long as heat is applied. Bumping becomes stronger when all salt is mixed with liquid.

There is almost no plating activity. A very small amount of silvery metal can be seen to form on the tip of the graphite cathode, but it doesn't thicken.

Conductivity is very low ( less than 2 [mA] current at >12 [V]. )

After a day and a half, I replaced 1CC of lost ethanol with 1CC of 1,3 propanol to see if solubility might be better in propanol.

Solution immediately darkens to a brown color, and conductivity doubles. But solution still appears to be colloidal.

I used an inkbird to temperature regulate the test tube over night. There are some risks as it disconnects occasionally (Randomly once in 12 hours, but sometimes it runs for 36 hours straight and reliably). eg: I had to program the heating unit to shut off the iron whenever temperature monitoring stalled for more than 60 seconds.

Electrolysis does release small amounts of hydrogen gas, but there was no sulfur dioxide smell. But, conductivity did not rise (significantly) as water was removed.


First picture, 102 [°C] at bottom of tube, roughly 75[°C] where graphite electrode is.
vlcsnap-2025-12-12-16h41m40s542.png - 191kB

Second picture, slightly different lighting, same tube with a few milli amps of current flowing through 6 washers. Hydrogen/oxygen bubbles are visible.


vlcsnap-2025-12-12-16h51m50s140.png - 239kB

After scratching off the graphite electrode, it immediately plated again with a thin grey coating.

I'm running AC current for a while to see if I can get more ions into solution while removing hydrogen and oxygen...

The solution is slowly becoming less colloidal and more of a clear brown liquid.
The colloid precipitated onto the washers in the background, and you can see some of it piled up on the side of the graphite electrode.

Parts of the graphite electrode, which were not scratched clean, did not plate after re-inserting into the solution. Mostly on the left side near the tip you can see darker color material.


[Edited on 13-12-2025 by semiconductive]

semiconductive - 13-12-2025 at 11:53

Hmm .... I see sodium dithionite also exists. Wow.
That's even more unstable than meta-bisulfite.

In all cases, the presence of at least one sodium cation is responsible for keeping the sulfer dioxide in solution for short periods of time before decomposition happens.

I don't see why lithium or potassium woudn't do the same thing.

If I've understood what I've read, correctly:

Sodium meta-bisulfite, dis-associates into two sodium-hydrogen-sulfites; (half salts), in water. That is equivalent to removing half the sodium from sodium sulfite while in solution. So the half sulfite salt must be reasonably stable in water solution. ( slow decomposition ).

If I want to replace the sodium and bisulfite ions with lithium or potassium bisulfite ions in alcohol, I must work out solubilities and ionization constants for a metastasis reactions to figure out which ions will exchange, and which ones won't. ( I still need to figure out the formulas for auto-ionization of alcohol. )

But it brings up two thoughts:

Perhaps I can electrolytically dis-associate iron pyrite (FeS₂ )into a solution of sodium metabisulfite. Iron can be in the +2 or +3 oxidation state. Sulfur can act like oxygen, S⁻², Therefore: I think iron pyrite might dissolve (on average) into solution as FeS₂⁻ ions.

But, if that's the case, then the ions might come from the cathode and not from the anode ?!

Ahh.... this might explain an earlier experiment that I couldn't reproduce. When I put aluminum anodes into an iron pyrite powder bath, and pulsed large amounts of current, sometimes I would get iron pyrite films rapidly forming on the aluminum anode surface. But, it wasn't consistent.

Note: Found a useful article
https://ajsonline.org/article/59780.pdf

FeS₂⁻ is unlikely in water, and I suspect alcohol:

Apparently, the most likely situation is ferrous ions and polysulfide anions in a hot water solution 40 [°C]:

Fe⁺² + S₅S⁻² + H₂O + HS⁻ → FeS₂ + S₄S⁻ + H₃O⁺

Sodium sulfide + solid sulfur + alcohol, might make a decent electrolyte to try.

Second thought:

I have cellulose acetate which I cam make a semi-permeable membrane that will allow positive ions to pass in alcohol -- but it will block negative ions. ( I can't use acetone with cellulose acetate, though, only alcohol or water -- because it dissolves in acetone. )

But:
If I put medium amounts of sodium metabisulfite in one compartment along with iron (solid) next to another compartment containing only sulfamic acid; (all materials submerged in ethanol and/or 1,3-propanol and separated by cellulose acetate) ,

I imagine the sodium atom will work it's way through the semi-permeable membrane fairly easily and can be precipitated out as sodium sulfamate.

Perhaps this would allow me to build up iron sulfide ions in solution on the other side of the membrane ?

(Any thoughts?)


[Edited on 14-12-2025 by semiconductive]

DraconicAcid - 13-12-2025 at 12:25

I think you'll have a hard time finding something that iron pyrite will be soluble in.

semiconductive - 13-12-2025 at 15:17

Quote:
I think you'll have a hard time finding something that iron pyrite will be soluble in.


Yes. So far, I have only gotten small amounts to transfer when using citric acid and DMSO.
But, it's only thin films formed repeatably. Thick plating only happens randomly.

But, the article in the previous post might explain what's going on.
Solid sulfur might need to be present to make poly-sulfides in solution.

Unfortunately, either I don't understand the author's notation or the reactions are not entirely balanced. I wrote into the post what I think they meant. ( Correct me if I'm wrong. )

I either need to add sulfur to the mix, or remove some iron from the pyrite in order that excess sulfur exists.

Hmm...
I think sulfur goes liquid at around 120[°C].
If it doesn't burn ethyl citrate at that temperature ... I can get my test tube that hot under kerosene, just fine....



[Edited on 14-12-2025 by semiconductive]

semiconductive - 13-12-2025 at 22:02

Now to figure out ionization constants...

I'll take the Kw data for ultra-pure water that I linked to earlier, and I'll assume basic semi-conductor physics (for ice like substances).

I looked up codata for kBoltzman = 8.61733326·10⁻⁵ [eV/K]
So. this will be more accurate than my earlier post's approximations.

The equation for the energy gap is:

C₀ = unknown and is affected by material compression/etc.
C₁ = 3/(2·ln(10)) ≈ .65144
C₂ = 1/( 2·kB·ln(10) ) ≈ 2519.889

E in terms of ( Kw [negative exponent] ,T [Kelvin] ):
E = ( Kw - C₀ - C₁·ln(T) )·T/C₂

At 25 [°C] = 298.15 [K], and 'ultra pure water' I know:
Eg = -0.8280 = ( -13.9933 - C₀ - C₁·ln( 298.15 ) ) · 298.15 / C₂
At standard lab conditions, 298.15 [K] or 25 [°C]:
C₀ ≈ -10.7069

Therefore, I get the following values using a linear energy gap correction that is *opposite* of what NIST shows.

E ≈ -0.8280 + .0008360 · ( T-298.15 )

This is from 0[°C] to 100[°C] in 5 degree increments:
Note: Exponents are negative, and I'm keeping the sign I calculated. For reporting pH calculations in Chemistry, the sign needs to be reversed.

Calculated:
['-14.88', '-14.69', '-14.51', '-14.33', '-14.16', '-13.99', '-13.83', '-13.68', '-13.53', '-13.38', '-13.24', '-13.10', '-12.96', '-12.83', '-12.71', '-12.58', '-12.47', '-12.35', '-12.24', '-12.12', '-12.02']

From the earlier linked reference, but rounded off to two digits after the decimal:
['-14.94', '-14.73', '-14.53', '-14.34', '-14.16', '-13.99', '-13.83', '-13.68', '-13.54', '-13.40', '-13.27', '-13.15', '-13.03', '-12.92', '-12.81', '-12.70', '-12.61', '-12.52', '-12.43', '-12.34', '-12.26']

Therefore:
With no phase change correction and assuming a linear model whose voltage *decreases* in total magnitude with increasing temperature; my basic semiconductor equation yields water errors of:

| 14.88 - 14.94 |/14.94 to | 12.02-12.26 |/12.26 = 0.4% to 2.0%.

(No surprise) The error is smaller near freezing (solid-state) than boiling.

The basic semiconductor equation, even without a correction for liquid motion or pressure vessel distortions is surprisingly accurate when I use the wrong sign of voltage change from the NIST paper because it agrees with my intuition (I'm not a chemist!). I expect semi-conductor band gaps to decrease in magnitude with increasing temperature...

Next: I'll compute a correction for typical material expansion and packaging in semiconductors, and see if I can get a better fit.


[Edited on 14-12-2025 by semiconductive]

semiconductive - 15-12-2025 at 00:21

End of control experiment, and it's not really good news:

AC current run for 24+ hours has not increased ion-conductivity of the solution.
I put a fresh graphite catholde (-) into the solution.

No electroplating activity is seen except a tiny bit near the tip. But the amount of metal is small enough that it might be an impurity and not necessarily iron.

The solution has become clearer and less brown with time.
I raised the temperature at the test tube bottom for the last 8 hours to 120[°C], and that just accelerated the clarification of the solution. The ethyl citrate is stable, no burning, and surprisingly I haven't lost another CC of solution by boiling out more ethanol.

Post mortem:
H₂SO₄ Ka1 = 100%, Ka2 = 1.2·10⁻²
Citrate Ka1 = 7.4·10⁻⁴ Ka2 = 1.7·10⁻⁵ Ka3 = 4.0·10⁻⁷ # zero ionic strength

Hmmm.... I don't recall how ioninc strength of organic acids change with concentration. But if I look at the zero ioninc strength, it suggests my mistake was thinking that the Ka2 of sulfate would trap a signifiant portion of the Ka1 from citrate.

The organic acids are much better at holding on to the lithium at low concentrations than the sulfate is. That's rather counter-intuitive. ( Who knows what temperatures did to the values... )

If I've only got around 10⁻² difference in Ka values at room temperature, then I suspect less than 1% of the lithium ions would transfer from citrate to sulfate at room termperature?

If I were to repeat the experiment with sulfite ions, though, the sulfite is a much weaker acid than sulfuric. It will hold onto sodium, lithium, and iron better. Unfortunately, I've already got sodium attached to it....

I'll finish my calculations for auto-ioninzation of water, then alcohol, and then try to work out some estimates for the same experiment using sodium meta-bisulfite rather than sulfate.






semiconductive - 15-12-2025 at 15:07

Rereading the pyrite article, both sides need to have negative 1 charge total:
I think I missed a '2' for the poly-suflide ion.

Fe⁺² + S₅S⁻² + H₂O + HS⁻ → FeS₂ + S₄S⁻² + H₃O⁺


If both sides are charge neutral, than possibly there is a missing/implied hydronium ion to neutralize the ionised HS⁻:

Fe⁺² + S₅S⁻² + H₂O + HS⁻ + H₃O⁺ → FeS₂ + S₄S⁻² + 2·H₃O⁺

I'm thinking: a very similar reaction might be possible if I use sodium sulfide salt: Na₂S and cook it with elemental sulfur, to produce polysulfide ions, in an alcohol solution.


I am able to find analogous reactions using either selenide or arsenide, so perhaps practical information and clues are available from more popular research

https://www.mdpi.com/2079-6412/13/11/1905


[Edited on 16-12-2025 by semiconductive]

semiconductive - 15-12-2025 at 22:50

Sigh. There is definitely a difference between semiconductor physics and chemical reactions / fuel cells. I find it amazing that I got such close agreement (2%) after curve fitting semiconductor equations when they are likely incompatible.

When I calculate the energy gap voltage for 'super pure water' using semiconductor equations, the energy gap can be shown to decrease as temperature increases.

Assuming the energy gap was 0.828 [eV] at 25 [°C], this graph shows the energy gap required for super pure water vs. temperature in order to produce the correct number of hydronium and hydroxide ions.

pngsnap.png - 9kB

Y axis is the energy gap, x-axis is temperature in Celsius.
This is a graph of the expected (empirical) band gap voltage shape vs. temperature is water acted the same as a solid semiconductor.

Note:
Although the ionizing voltage drops as temperature increases, the ionizing voltage changes less and curves more with increasing temperature.

That's a clue that the physics is very different.
Y.P. Varshni's correction isn't going to work for water.

The trend in semiconductors is opposite of what water does. A typical semiconductor ionization voltage curves most near absolute zero Kelvin, and decreases in magnitude with increasing temperature -- but the slope of energy gap change per decree celsius becomes more linear with temperature (not less).

example: Silicon, and several other semiconductors:
https://www.researchgate.net/publication/319068163_A_novel_t...

When I read up on fuel cell reactions, I see that the voltages measured do increase with temperature. ( Just tried it with a AA battery as well. )

But when I measure the voltage across a silicon diode for a fixed amount of current, the opposite happens. As the diode gets hotter, the voltage decreases.

So my intuition is exactly backward, and I need to figure out why before I can do anything more.








bnull - 16-12-2025 at 08:22

Quote:
There is definitely a difference between semiconductor physics and chemical reactions / fuel cells. I find it amazing that I got such close agreement (2%) after curve fitting semiconductor equations when they are likely incompatible.

I've had my share of these things in my time in Physics. After a while, the amazement gave way to a chuckle and a "That again."

Some comments. (1) pH decreases with temperature, and pH plus pOH is not a constant. The sum is 14 at 20 °C (or 25 °C, I forgot which one) and goes up or down according to how much hotter or colder than that water happens to be. (2) The lattice in solid water is very different from the one in semiconductors. Water molecules are polarized, whether protons or hydroxyls are present as impurities or not. The same doesn't happen to silicon, not to mention that the impurities in silicon serve to increase conductivity. The ways that charges can travel within both lattices are very different. It's years since I dealt with semiconductor physics and I forgot most of it. I can still visualize it but I can't explain it in words.

Edit: Fixed an idiotic mix up.

[Edited on 16-12-2025 by bnull]

semiconductive - 16-12-2025 at 10:22

Quote:
Some comments. (1) pH decreases with temperature, and pH plus pOH is not a constant. The sum is 14 at 20 °C (or 25 °C, I forgot which one) and goes up or down according to how much colder or hotter than that water happens to be.


Charge neutrality must exist both in intrinsic semiconductors and (equally true) neutral liquids. Electrons and protons are neither created nor destroyed during ionization events -- the charges only physically move around.

For: Pure water (with no contaminating ions that are not made of hydrogen and hydroxide); I already know the Kw data. SInce pOH=pH at every temperature that is electrically charge neutral -- I expect pOH+pH = 2·pH = 2·pOH for distilled water.

From two posts back the exponents of Kw for 'ultra pure' water are listed. I think these values equal -(pH+pOH) for every 5 [°C] increment:

['-14.94', '-14.73', '-14.53', '-14.34', '-14.16', '-13.99', '-13.83', '-13.68', '-13.54', '-13.40', '-13.27', '-13.15', '-13.03', '-12.92', '-12.81', '-12.70', '-12.61', '-12.52', '-12.43', '-12.34', '-12.26']

Therefore, the pH value of 'neutral' water is (by definition) - 1/2 the total exponent:

pH=pOH=[' 7.47', ' 7.36', ' 7.26', ' 7.17', ' 7.08', ' 7.00', ' 6.92', ' 6.84', ' 6.77', ' 6.70', ' 6.64', ' 6.57', ' 6.52', ' 6.46', ' 6.41', ' 6.35', ' 6.31', ' 6.26', ' 6.21', ' 6.17', ' 6.13']

T[°C] = [ 0,5,10,15,20,25,30,35,40,45,50,55,60,65,70,75,80,85,90,95,100]

The number 7.00 shows up at 25 [°C] in this chart.

Therefore: I'm seeing distilled water pH values decrease with temperature.

eg: That means the number of hydronium ions is *increasing* as the liquid gets hotter because the number's exponent is by convention the negative of the pH number.

Both authors I found online show the same trends (though slightly different values) vs. temperature. They agree very well from 0 to 30 [°C], but there are disputes up to 3% for hotter temperatures.

Note: There are a couple of problems with this kind of data: people regularly fail to report what the source of water was and how the measurement was made.

Therefore, I can't do any math to correct for things such as an experiment done in a closed jar (approximately iso-choric) vs. an open jar (approximately iso-baric). In a lot of ways, this makes the data somewhat useless...!

But: Two different authors have the given approximately the same values (rounded to three digits) in their data for low temperatures, so I think whatever experimental conditions were used by one author are very similar to the other author.

I could put some pH indicator in distilled water, but that technically will change the pH since pH is very sensitive to the mass-action law.

Note: The same mass-action law is used in semiconductors as with liquids.


[Edited on 16-12-2025 by semiconductive]

bnull - 16-12-2025 at 10:47

Quote:
I'm curious: Do you have a reference experiment showing the pH increase with temperature?

No. I just happen to be a jackass and didn't notice what autocorrect did and wrote accordingly. Sorry for that.

Edit: What I was going to write, and somehow fumble and forgot it, was that the number of charge carriers increase with temperature, and so does the movement of the water molecules. In semiconductors, such movement is restricted by the lattice. The same doesn't happen in water because there is no restriction as to where a molecule goes except for the walls of the container and (to some extent) the surface of the liquid.

Also, the band gap is directly related to the characteristics of the lattice; what happens to the band gap when there is no lattice?

On a side note, ice can behave as a semiconductor. I've downloaded a paper (which I intend to read as soon as I can find it) about electrical/electronic/electrochemical properties of ice, with and without dopants (mainly acids and bases). This, unfortunately, is completely useless for you as you want to plate metals onto stuff, not make an ice transistor.

[Edited on 16-12-2025 by bnull]

semiconductive - 16-12-2025 at 14:07

Quote:
Edit: What I was going to write, and somehow fumble and forgot it, was that the number of charge carriers increase with temperature, and so does the movement of the water molecules. In semiconductors, such movement is restricted by the lattice. The same doesn't happen in water because there is no restriction as to where a molecule goes except for the walls of the container and (to some extent) the surface of the liquid.

Also, the band gap is directly related to the characteristics of the lattice; what happens to the band gap when there is no lattice?


OK! Got it!

Hey, every post of mine is edited ... I'm lucky to write exactly what I mean by the third edit. ( Disability and medication side-effects. )

I'm very qualified to do solid state semiconductor modeling. So good that a patent of mine has been stolen by 13+ U.S. companies who never paid the owner of the patent a penny. Starting cost for litigation is $6 Million. ( And I'm not bothering to pursue. )

So, I'm kind of reading your question and thinking "He believes the grass is greener on the other side of the fence AKA: life is simpler in solids."

I agree, there are differences between liquids and solids; and I'm going to have to understand them to get any better at chemistry.
But, there is a reason I tried the semiconductor equations on liquids; and I'd like to elucidate a bit.


AKA: solids aren't simple! (Trust me!!!)

If you click on the link I gave for energy gaps in semiconductors, you'll see both simple crystalline semiconductors and compound semiconductors with huge mixes of oddly shaped atoms (both). The crystal shape (no matter how perverted) doesn't change the basic equations which model them.

The fact is, many solids even have dipole moments in the crystal. A good example is the lithium citriate I made for the test experiment in this thread. As a solid, Cit-Lit is piezoelectric; therefore merely placing an electric field across a crystal will cause the lattice spacing to change. ( eg: for the crystal to snap meta-stably into different shapes. ) This is how "Ferroelectric" memory works.

The 'lattice spacing' of silicon isn't really constant, either.
Depending on which direction you go measure the atom spacing through the crystal, the atoms will be closer or farther apart.

I can even take electronic transistors in plastic packages, and make a device to change the gain of a purchased transistor changing the pressure applied to the outside of the package.
Very few people even realize that merely talking next to a transistor can cause it's gain to change a *tiny* bit from sound waves hitting it.

The ability of lattice spacing (in silicon/solids) to change just means that the tendency of ions to be released in certain directions, and the conductivity of the material in different directions, is hetrogenous.

eg:
I can demonstrate metal contacts hooked to silicon wafers at exactly the same spacing, but along different directions of the crystal, and the resistance values measured by my ohm meter will be different in one of three orientations (but not the other two).

Anytime you put an impurity atom into a lattice, the lattice gets deformed. Other times, ions can migrate through the lattice just as if it were very viscous liquid. Electromigration is a real phenomena even in AMD and Intel made microprocessors.
Their engineers do everything they can to 'stop' it!!!!

I seem to recall; You linked me in another thread to an article about making sodium by having it flow through the glass envelope of a vacuum tube. The electrons from a heated filament reduce the ions to sodium metal.

So, I'll return your own novel example to you.

That, right there, is solid state electroplating!
( Cool idea, by the way. )

Electroplating (albeit very SLOW) is also likely possible in ice.

It's also the case that semiconductors can melt when they get hot enough. Yet, (strangely) the same equations are used to model them.

So -- with that backgorund: let me answer your rhetorical question:

There always an average distance between atoms, and you do a density of states calculation in all directions and compute a statistical average 'effective' lattice distance that weights the distances by how often (percentage) they exist.

When you get into liquids, the same ought to work.

I think the article that Draconic Acid linked me to demonstrates this for alcohol:
Quote:

For methanol, pK(autoionization) = 22.67

https://pubs.acs.org/doi/10.1021/acs.jpca.5c03979

The quantum mechanics relationships are identical regardless of the phase of the material for density of states calculations.

I'm thinking:
The very fact that the law of mass action is used in chemistry of liquids and that the same law can be derived physically from density of states calculations in solids (semiconductors), shows that the fundamental physics can not be essentially different; but only that the density of states needs to be modified in some way in *value* or trends.

I'm sure if you get a chart of the density of water (molecules/cm³) from NIST, it will show that water molecule spacing (average) expands over most of it's liquid temperature range from 5[°C] to boiling.

When I thought about doing the calculation, I only suspected a real difference in water properties as compared to silicon in the -10 degrees to +10 degrees celsius region.

Because this temperature range is a place where the spacing of atoms doesn't follow a single trend. It doesn't always shrink with temperature or always grow with temperature.

Y.P. Varshni's correction assumes a single direction of change.

I'm sure theres a couple of differences in that:

1) The spacing of molecules can change rapidly with time in a non-oscillatory (phonon) manner in a liquid. 2) The molecules are able to rotate which allows for more 'states' to exist.

Thus the 'density' of states will be affected by how much a molecule is free to rotate vs. temperature.

But the fact that the calculation reveals the band gap becoming almost constant near the boiling temperature of water means that there are at least two spacing/fighting effects going on in water that work in opposite directions.

From the chart, it looks like 40 degrees C is about where these two different trends in statistical spacing of water happens.

I don't see anything suspicious near 4 degrees Celsius, and that's where I would have suspected the biggest differences to appear. The density of water changes it's slope around 4 [°C].

This makes me suspicious that I've mis-identified the band-gap (or band separation) energy of water ionization with the voltage of the fuel cell.

In semiconductors, I used to laugh at people who think the optical band gap is identical with the semiconductor band gap. For, I've done experiments showing they don't always agree -- but that there is a predictable relationship in the error.

But, I've got that feeling that I'm probably making the same mistake in this thread and I'm not sure where I did it. Anxiety kicks in....

Boltzmann statistics and math operate exactly the same regardless of material phase. So, what assumption have I made about batteries that is wrong -- and how do I check it?


bnull - 16-12-2025 at 14:43

My statistical mechanics has been dead and buried since before Pandemics. Good to see forgotten stuff brought back to life (of sorts). And no, I've no answer to that.

I found the article I mentioned and there's good news and bad news. The good news is that the article on semiconducting ice mentions another article on the measurement of the band gap of water, which by its turn references an article on the dependence of the ionic product of water (Kw) on temperature (in Russian, again), so you can compare the latter with what you have so far. The bad news is, the publisher of both the first and the second article is considered a predatory publisher and they were written by the same person. But even a broken clock is right twice a day, so I'm attaching both articles here.

Attachment: S. Yefimov - Ice diode.pdf (378kB)
This file has been downloaded 35 times

Attachment: S. Yefimov - Band gap of water.pdf (292kB)
This file has been downloaded 38 times

semiconductive - 16-12-2025 at 17:14

Very interesting. Thank you for the articles.

Note: 4.3[°C] is roughly 0.0036 on the chart in the second paper.

S. Yefimov is assuming a constant energy gap for all data points and using a simplified equation.

If I plot his paper's regression line vs. the combined (rounded off) American Kw data that I found, you'll see that my plot and his are very similar except at 4.3 degrees C.

( Which is where I would expect a defective point to be found based on water density being a maximum. )

Note: When I plot his equation (2), I don't get his regression line.
Do you see a mistake in the equation I plotted?

pngsnap.png - 12kB

I will recompute the band-gap value from the regression line of the plot (in brown), to see what I get (next post) if there's no obvious error in what I plotted.

semiconductive - 16-12-2025 at 20:16

There's only one really noticeable difference between the equations S. Yefimov is using, and the ones I used from quantum mechanics.

He isn't correcting for quantization of states.
eg: He's neglecting a T^(3/2) factor which is missing.

Why:
Whenever an object moves it has an energy and a momentum; but quantum mechanics makes a single extra requirement that all possible energies and momentum(s) are not continuous but are broken up by Plank's constant and the relationship E=hf.

What this ultimately means is that there is a discrete number of speeds which an object of mass 'm' can take on for finite changes in energy. The number of speeds an object can take on for a finite change in energy is known as the 'density of states' (DOS).

The following videos are not necessary, but they are a refresher course on how to figure out how many charged particles of a given mass could be moving (in any way, whatsoever) in a trapped (ionized) environment.

https://www.youtube.com/watch?v=z7YGS67GETo
https://www.youtube.com/watch?v=3vFNQOx6kBo

To summarize the video: for charged objects having some average mass, the statistical number of possible ions per mole of ionizable material (where the ionized object has mass m) is determined by a simple classical approximation formula:

g(E)·dE = 4·π·(2·m)^(3/2)·√(E) / h³

This formula is written in the square root of energy, but it really is linear in velocity (speed of ion travel.) And chemists know that velocity is fixed by temperature and mass.

Therefore, the full equation for semiconducor ion concentration is usually a little bit more complicated than in the articles you linked.

Ni = C₁ · T^(3/2) · e^(- Eg/( 2·k·T ))

Consider:
The number of atoms in a given mass of water, or silicon, or whatever, is usually constant. But the volume that number of atoms occupies changes with temperature depending on the thermal expansion characteristic of the material.

Hence, the above formula should (personal opinion) work even for gasses, but it doesn't take into account volumetric changes.

In the other article:
The straight line fit, 2670.343/T + 5.036 assumes that the volume of water does not change with energy and the water does not change in volume with temperature. eg: He's assuming these two effects cancel each-other out -- and they generally don't.

It's the gentle curve among the data points which he is not predicting, correctly.

In my formula, you can remove the term which makes a generic quantum mechanical correction for ion velocity. This will reduce it to the same as S. Yefimov's equation.

Doing that means:

Eg = (kW - const )·T / 2519.8890927
(kW-const) = Eg*2519.8890927/T

Assuming the constant is zero (as I don't know why the constant shows up there, anyhow), the energy gap in electron volts for the article ought to be:

2670.343/2519.8890927 = 1.0597 [eV]
























[Edited on 17-12-2025 by semiconductive]

chornedsnorkack - 16-12-2025 at 23:44

What is the distinction between "electromigration", "electrolysis" and "electrophoresis"?

bnull - 17-12-2025 at 03:11

Quote: Originally posted by chornedsnorkack  
What is the distinction between "electromigration", "electrolysis" and "electrophoresis"?

Electromigration: https://en.wikipedia.org/wiki/Electromigration; electrolysis: https://en.wikipedia.org/wiki/Electrolysis; electrophoresis: https://en.wikipedia.org/wiki/Electrophoresis. It can't get much better than that, there's no confusion or controversy among the terms.

Or, if you prefer a short and rather incomplete version: electromigration is when atoms in a conductor move along the material due to transfer of momentum from electrons to these atoms; electrolysis is the decomposition of an electrolyte by means of an electric current; and electrophoresis is a technique to separate (polar) molecules by the application of an electric field. Notice that electromigration requires the physical interaction between electrons and atoms, while electrophoresis uses an electric field.

Edit: Fixed it.

[Edited on 17-12-2025 by bnull]

semiconductive - 17-12-2025 at 14:31

Quote: Originally posted by bnull  

Or, if you prefer a short and rather incomplete version: electromigration is when atoms in a conductor move along the material due to transfer of momentum from electrons to these atoms; electrolysis is the decomposition of an electrolyte by means of an electric current; and electrophoresis is a technique to separate (polar) molecules by the application of an electric field. Notice that electromigration requires the physical interaction between electrons and atoms, while electrophoresis uses an electric field.


Hmm.
Quote:
On a side note, ice can behave as a semiconductor. I've downloaded a paper (which I intend to read as soon as I can find it) about electrical/electronic/electrochemical properties of ice, with and without dopants (mainly acids and bases). This, unfortunately, is completely useless for you as you want to plate metals onto stuff, not make an ice transistor.


Combining your quotes:

So when the wire representing 'move direction south, by southwest' really fast in the Russian RS28 SARMAT missile suddenly found itself on the 'math coprocessor emergency interrupt' -- that's an example of electromigration.

But, it's probably not considered electroplating unless the wire moves through silicon in order to find itself on top of the math co-processor?

---
If you prefer a rather (shorter) and incomplete answer to your quest to understand 'What happens when there isn't a crystal lattice?'

One answer was hidden very cleverly in the "statistical mechanics" paper I already linked to, on page 17. ( First sentences of the left columb. )

https://srd.nist.gov/JPCRD/jpcrd696.pdf


What "you do" is invent a new word called "quasilattice".
This word, of course, means that you go right ahead and apply the equations meant for crystals to steam vapor.

It's a bit like when my physics professor said, "Now suppose we use a spherical model for a cat."...

bnull - 17-12-2025 at 15:12

The first quote refers to @chornedsnorkack's question that unfortunately ended up at the end of the previous page, making my pointing out the definitions a non sequitur.
Quote: Originally posted by chornedsnorkack  
What is the distinction between "electromigration", "electrolysis" and "electrophoresis"?


It wasn't a quest. It was a passing question, like those comments in Fieser and Fieser's Organic Chemistry. "Oxalic acid can be used to dehydrate cyclohexanol," that sort of thing. But it's good to know what to do in such cases.

As far as I know, the definition of electroplating involves a surface and a medium in contact with the surface from which comes the plating material that is deposited in the surface. The plating material traveling through the body whose surface is to be plated doesn't seem to fit the definition. Maybe one could call it electromigration plating. Theoretically interesting (yes, it is) but probably useless when one wants to plate stuff.

[Edited on 17-12-2025 by bnull]

semiconductive - 17-12-2025 at 17:11

All I said is taken in good humor, I hope.

There's a lot of things that I don't mention, because I write too much for most people to read (as it is).

You might not know that when the graphite electrode of my last experiment touches the glass wall of the test tube that gas bubbles erupt much faster than if it is not touching the glass.

This odd detail made me consider your article and dictionary definitions in ways I'm not going to fully explain. But: In the surprisingly novel experiment you posted there are several things potentially going on, that I'm not even sure how I would talk about them (vocabulary wise).

One example: Do the charge 'carriers' really move through the ice in his experiment, or do they travel along the surface of the copper wire to the surface of the ice and then migrate toward the 'junction' ?

( "Obviously", Ice melts along its surfaces ... and maybe under pressure at junctions. )
It's not just electrons which can move along ice surfaces.

But: The equation I posted for density of states (DOS) was derived presuming the only thing that actually moves are electrons (and spaces left where electrons ought to have been -- holes);

Therefore the only units of energy needed is 'electron volts'. But, if I did the same (DOS) derivation assuming protons as charge carriers, then the final equation would have to be adjusted to have different valued constants and even the exponent might change. ( There are no proton-volts units, you just have to multiply electron volts by some scaling factor. And, electrons spin in pairs ... but is this really relevant to protons? )

Final passing comments:
I've not been talking about Einstein's E=mc², but measuring relative permittivity and not getting the value '1' is equivalent to saying the value of 'c' is different inside materials than in empty space.

This means the relationship between Energy and Mass ( which are the only two things used in the Density of States formula derivation) are very slightly different inside a material than outside of it.

Now: I'm thinking --

The major difference between liquid and gas -- is that gasses tend toward a constant number of moles of material per volume; P·V=nRT (ideally).

In a constant pressure situation, P is not allowed to change. But that means the density of the substance must change drastically at the boiling point -- and therefore, so does the dielectric constant of water and steam.

See my plot of energy gap back a few posts?

I think the rapid change of energy gap shape near 100 [°C] in my plot isn't a math error (in spite of the specific value I chose from NIST data at 25C being likely wrong).

It seems a reasonable hypothesis that normal water boils at a slightly hotter temperature than ultra pure water. eg: The strong (and unusual) curve bend of the Energy gap is likely evidence that the experiment was either done slightly above sea level, or that ultra pure water boils at a lower temperature than normal distilled water.

semiconductive - 17-12-2025 at 18:02

I see my plotting mistake! I typed in 0.03033. and not 0.003033 into my plot.
This means S. Yefimov's equation as written does produce the same regression line.

But why do I get an estimate of >1 electron volt for his line, when he gets a value of 0.53 electron volts....

"A broken clock is correct twice a day"... :)

Maybe:
0.53 [eV] · 2 = 1.06 [eV] which is rounded off 1.0597 [eV] of my number.

--- For future reference ---
IUPAC defines the boiling point of water under slightly different conditions.
Does Russia generally follow IUPAC?
I wonder if I will need to do the same for alcohol.... sigh...

96485.3321233 [eV] → 1 [kJ/mol]

Melting H₂O = 334 [ KJ/Kg ] ≈ 6.01701 [k J/mol ] ≈ 6.236 [ μ eV ]
Vaporizing H₂O (99.61 [°C] at 100 [k Pa] ) = 2257.5 [ k J/k g ] = 40.6688625 [k J/mol ] ≈ 42.15 [ μ eV ]

NOTE: Interesting discovery: ( I never knew this before. )

The amount of energy required to melt ice or to boil a single molecule of water is so small compared to the ionization bandgap of water itself, that energy discontinuities during phase changes will not show up in a semiconductor band-gap plot.

AKA: It's not practical to be able to detect a 40 micro-volt change with a desktop volt-meter, reliably, or to show it in a plot.

This means that it it physical volume changes that are messing up my band-gap plot, and not electronic changes of individual molecules of water.

Kw ionization data is (apparently) very sensitive to dissolved gasses in a fluid (micro-bubbles) and even NON-ionizable substances in contact with the fluid such as plastic container walls, dissolved droplets of kerosene or oil, etc.

To make an accurate band-gap estimate for semiconducting water, the KW data needs to be divided by the actual density changes of the water itself (isolated) from other density changes.

eg: There can be no accurate water density value at 100 [°C] as that could be either water or steam, (depending on experimental setup , and time given after temperature change to 'equilibriate'.)

It's not reasonable to believe these experiments were carried out in the international space station with a heater at the center -- and that means the pressure of the fluid must not be perfectly constant, but rather pressure must form a gradient that is lowest near the water-gas interface.

Therefore:
For accuracy and precision reasons: I need to discard Kw data at 100 [°C]. The remaining Kw data needs to be adjusted for density of water changes before it can be used to compute semiconducting coefficient values.

https://nvlpubs.nist.gov/nistpubs/jres/097/jresv97n3p335_a1b...


[Edited on 18-12-2025 by semiconductive]

semiconductive - 19-12-2025 at 10:43

NIST already has water density equations that are accurate from 5 [°C] to 40 [°C], for degased or gassy water.

It's difficult to find water ice data where I know whether the water was boiled before freezing or not. But, I can find single crystal experiments which will be immune to dissolved gas problems.

Example: "Thermal Expansion of Single-Crystal H2O and D2O Ice Ih", Physical review letters 121.

All articles I can find show that hexagonal H₂O Ice hits it's maximum density at around 60 to 64 [K]. The graphs also agree that on average hexagonal ice expansion is isotropic.

I can fit the H₂O graphs in the cited article (by eye) with the equation:

DL/L₁₀ ≈ 138.2·10⁻⁹·(T-62)² -0.135·10⁻³

Scaling this equation linearly and then cubing to compute a constant mass ice volume fails with wrong values. This is not unusual -- https://www.physicsforums.com/threads/calculating-the-coeffi...

Alternate attack, I get two data points for ice density online.
T=-20 [°C] = 273.15 [K] → 0.9196 [ g/cm² ]
T=0 [°C] = 273.15 [K] → 0.9167 [ g/cm² ]

What I want is an equation that hits these two data points and has a maximum density at 62 Kelvin.
Density ≈ 1/( a·( T-62 )² + b )³

From 38.15 [K] to 273.15 [K], I can approximate:
Density=1/( 134.631·10⁻⁹·( T-62 )² + 1.02341 )³ [ g/cm³ ]

From 38.15 to 273.15 every 5 degrees:
['0.93272', '0.93280', '0.93286', '0.93291', '0.93293', '0.93293', '0.93292', '0.93289', '0.93284', '0.93277', '0.93268', '0.93258', '0.93245', '0.93231', '0.93215', '0.93197', '0.93177', '0.93156', '0.93132', '0.93107', '0.93080', '0.93051', '0.93021', '0.92988', '0.92954', '0.92918', '0.92880', '0.92840', '0.92798', '0.92755', '0.92710', '0.92663', '0.92614', '0.92564', '0.92511', '0.92457', '0.92401', '0.92344', '0.92284', '0.92223', '0.92160', '0.92096', '0.92029', '0.91961', '0.91891', '0.91819', '0.91746', '0.91671']

For water near freezing, I can find data:
T=0⁺[°C] = 0.9998 [ g/cm² ]
T=3.98⁺[°C] = 1.0000 [ g/cm² ]

The zero degree value is actually the NIST ITS-90 polynomial rounded down, even though it's not technically valid between 0 and 5 [°C]. This is a reasonable approximation, so I'll keep it.

Composite density of H₂O chart from 38.15 [K] to 40 [°C], is then:

pngsnap.png - 19kB

For any given experiment, the steep discontinuity will not be present because the time for freezing will become confounded with the temperature.


[Edited on 19-12-2025 by semiconductive]

bnull - 19-12-2025 at 15:54

Quote:
All I said is taken in good humor, I hope.

Yes.

What software you're using to plot data?

semiconductive - 19-12-2025 at 16:34

I'm just using Gnuplot. Free, open software.

Note: I have a success!

If I adjust the densities of pH+pOH, to be Molal based rather than volume based, and don't force any 'p' values to match NIST, then the failure in bandgap characteristic goes away. I get an excellent semiconductor fit using Varshini's correction.

It doesn't matter if I use the Russian or American data set, the answer is approximately the same.

Less than 0.2% error, everywhere.
This is a reasonable semiconductor model of 2·H₂O ionization into H₃O⁺, OH⁻.
I'm computing the pH+pOH values vs. inverse temperature. ( 'x' axis is 1/Kelvins )

molal_phoh.png - 10kB

The maximum energy gap to overcome to ionize water and that fits all data points is 4674.12 / 2519.88... = 1.85 [eV]

This is not the usually published band-gap of water itself. Rather, this is difference in ion potentials between [OH⁻] and [H₃O⁺] which correspond to the conduction and valence bands of silicon. eg: this gap is the ionization energy to make two ions with the 'Fermi' level being halfway between.

The Varshini correction is small, which is a good sign that I've done the math right. The equation ought to be correct for both solid ice and water. (Possibly even steam, but that's tricky.)

For all data points, assuming NIST density equation is correct all the way to boiling.
pH+pOH (Molal) ≈ 4674.12/T -10.5102 +0.651441·ln(T) + 7.4969·10⁻³·T²/( T+1.36780)/log₁₀(e)

If I re-run the fit, using only data points from 0 to 40 [°C], the band gap maximum becomes 2.47 [eV] ; but the Varishini's correction becomes larger.
pH+pOH (Molal) ≈ 6232.50/T -21.0663 +0.651441·ln(T) + 15.2774·10⁻³·T²/( T+1.36996)/log₁₀(e)

This suggests that the NIST density equation likely is close to correct even above 40 [°C] because a smaller correction term usually indicates a better over-all physics match.

And now I'm a bit stuck. There's disagreement in data sets at high temperature (near boiling) of up to 3%. The QM correction I'm using is just the electron mass version (which exists even with real ions) -- but there is a more correct set of constants that reflect the relative permittivity of water ( As changed by E=m·c² due to 'c' having a smaller value in solvents of any kind. )

Still, even wrong, either equation is close enough to correct to do basic predictions with. Getting a more accurate equation requires experiments that I can't do yet.

I know that in semiconductors, when there are a lot of competing band transition values that depend on the direction you move through the crystal, the "band gap" of the material is always taken as the lowest ionization energy possible.

With water, the lowest ionization energy corresponds to physical hydroxide and hydronium ions. But, in the literature I find, they aren't reporting those ions energy differences as the band-gap of water.

But: There is a reasonable article I found, here, that tries to explain the different kinds of band gaps present in water.

https://www.researchgate.net/publication/276498338_Electroch...

But, I'm not sure where he's getting the K_H₂≈2·10⁻¹⁹ and K_O₂≈6·10⁻²² from.

Is that a Henry's law type of reasoning?

From: https://www.engineeringtoolbox.com/gases-solubility-water-d_...
where [c] is IUPAC symbol for Molarity, [ b ] is Molality:

I can see that at 25[°C], 1 atm.
[H₂]≈1.55·10⁻³[g]/2[g/mol] = 0.78·10⁻³ [c]
[O₂]≈0.04 [g]/31.996 [g/mol] = 1.2·10⁻³ [c]

I see a solubility of Oxygen that is 10x larger than he has in his paper.
Hmm...

reaction K₁: H₂ + 2·H₂O ⇌ 2·H₃O⁺ + 2·e⁻
K₁ = [ H₃O ]² / [ H₂ ] = [10⁻⁷]² / [ 0.78·10⁻³ ] = 1.3·10⁻⁹

reaction K₂: O₂ + 2·H₂O + 4·e⁻ ⇌ 4·OH⁻
K₂ = [ OH ]⁴ / [ O₂ ] = [10⁻⁷]⁴ / [ 1.2·10⁻³ ] = 8.3·10⁻²²

His oxygen constant value is close, but the Hydrogen value isn't.
I'm not sure what the reasoning is. :(



[Edited on 21-12-2025 by semiconductive]

semiconductive - 21-12-2025 at 10:06

Earlier, I tried to use the Nernst equation with fully balanced equation for an alkaline cell.
And I knew I was doing something wrong, but I couldn't figure out what.

2·[OH⁻] + H₂(g) ⟺ 2·H₂O(l) + 2e⁻
The voltage I looked up from NIST for the half cell reaction was: 0.828 [V] @ 25[°C]

https://www.sciencemadness.org/whisper/viewthread.php?tid=16...

Now, I've found an online tutoring company that happens to do the reaction:

https://allen.in/dn/qna/11044569

They list the full reaction as:
2·H₂O + 2e⁻ ⇌ 2·[OH] + H₂ -8.2777 [V] @ 25 [°C]

What I find fascinating about the video, is that it properly lists the full equation as the problem to do. But when you watch the tutor actually describe how the problem is done, she immediately changes the fully balanced reaction into a fractional equation with a different number of electrons. ( But doesn't explain WHY! )

H₂O + e⁻ ⇌ 1/2 H₂ + OH⁻ -0.8277 [V] @ 25 [°C]

So, obviously I forgot something in the 30+ years since I took chemistry in college. But, I still don't know what.

And then the tutor plugs the single electron version of the equation into Nernst.
I plugged in a two electron version.

At the end of her calculation, it's obvious that K she is calculates with the Nernst equation is actually Kw.

Now, I already have a semiconductor fit equation for Kw at *all* temperatures that is very accurate.
Therefore, without knowing why her problem works -- I can plug my Kw value into her Nernst equation (as written) and I should compute accurate full cell voltages for the reaction at any temperature. ( I'll do so in a lower post.)

I already suspect what is going to happen is that I will get cell voltages that increase with temperature, in spite of the fact that the band-gap for the reaction really is decreasing with temperature. ( I'll check in a post below. )

Note: I looked up the equilibrium constant for water spontaneously decomposing into non-ionized hydrogen and oxygen gas. ( Google's AI did it for me based on the full equation that I entered. Obviously, It could be wrong.... )

2·H₂O ⇌ 2·H₂ + O₂
K= 2·10⁻⁴² 25[°C] and 1 [bar]

But: I notice that the K value of this reaction is suspiciously close to the product of both K values in the paper I'm trying to figure out (immdiately preceeding post from this one.) The solubility of Oxygen was off by a factor of 10, this answer is also off by a factor of 10.

( K_H₂ · K_O₂ ) ≈ 2·10⁻¹⁹ · 6·10⁻²² = 1.2·10⁻⁴¹

Might just be a coincidence...

semiconductive - 22-12-2025 at 14:34

Next: I always carry out my calculations to more digits than really can be used. Then I throw away most of the work I've done.

Nernst Equation at 25 [°C] as used in the tutorial, but with extra useless digits:
E°_cell ≈ 0.0591593496847823 [eV] / 1 · log₁₀( Kw )

Oh, wait, the Nernst equation is only valid at 25[°C] ?!

Then what equation am I remembering from the newer library book I read at my alma-mater ?!

Kotz and Purcell, "Chemistry & chemical reactivity", (c)1987.
ISBN 0-03-058349-7

PP. 715 "The Nernst Equation"
I read ... blah blah ...

E°_cell = E° - R·T / ( n·F ) · ln( Q )

Oh, that's the Gibbs equation. It only gets re-named Nernst after someone plugs in T=298.15 [K] and computes the decimal value of the constants.
Q is any ratio of ions, whether in equilibrium or not. K is if they are in equilibrium.
Since we're doing auto-ionization, Q=K.

Hmm...
My equations are written in terms of log₁₀, and so is the final Nernst equation at 298.15 [°K] as stated in the book:

E°_cell = E° - R·T / ( n·F·ln(10) ) · log₁₀( K )

uh-oh. Maybe Q in my book can't be a function of temperature Q(T). They might mean a specific constant number Q° (at 25 [°C], 1 Bar. ); therefore K in the Gibbs equation, and the Nernst equation, could be K°. ( Bad words omitted. This is why I got the lowest grade in all my college work in Chemistry. They wrote K and not K° ?! )

Note: My semiconductor equation is already in the form of logarithms that vary with temperature: log₁₀( Kw(T) )

The simplified equation from Kotz and Purcell ought to be written:
E°_cell = E° - .00019842143... [V/K] · T/n · log₁₀( Kw° )

I was wrong, I can't plug my Kw into this formula and get the precise cell voltage at all temperatures. S***s to be me.

Kotz and Purcell list standard *reduction* potential on p. 710 for:
O₂(g) + 2·H₂O(l) + 4·e⁻ → 4·OH⁻(aq) as +0.4.... [V]
2·H⁺(aq) + 2·e⁻ → H₂(g) as 0.0.... [V]
2·H₂O(l) + 2·e⁻ → H₂(g) + 2·OH⁻(aq) as -0.827.. [V]

(Note: no Hydronium ion was mentioned in the table, just H⁺.)

Worse note:

No matter what Kw value (<1) that my equation computes for 25 [°C], the Gibbs equation as written in my chem book will make the cell voltage voltage become less negative with bigger temperatures. AKA: the log₁₀ value of ion concentration will *always* be negative and the gibbs equation multiplies it by another negative sign.

For situations where Kw < 1, and the cell magnitude is negative; the cell voltage magnitude must decrease toward zero according to the Gibbs equation in Kotz and Purcell.

But reality: I warm a battery in my hand, the voltage number/magnitude on my volt meter increases in value. This is true whether or not I reverse the probes on my meter and the voltage is artificially made negative or positive.

My present line of thinking:

The energy level difference between OH⁻ and H₃O⁺ (in electron volt energy units [eV] ) must decrease with temperature because of Boltzmann statistics and physics.
The gap must decrease with temperature in order to gently curve ionization concentrations non-linearly. eg: to make the gentle 'curve' in my plot.

If K is K° in the gibbs equation. Then the Gibbs equation and the Nernst equation are purely linear in temperature. This is *not* what people report the ionization constant of real water does with temperature.

But:
The equations shown in my book, and recorded in NIST have sign conventions that aren't explained well. In order for my book's equation to be correct, Kw must be >1. In that case, it doesn't make sense that my books equation is the same as the Nernst equation -- becuase the sign is NOT reversed when simplifying to disagree with the tutorial that I linked.

I know:
Water ionization is at around 1.2... [V] for each electron ionized in dilute water. In concentrated 40% aq HCl + titanium dioxide, Ive measured battery cell voltages as high as 2.6 [V] at room temperature...

But, I still don't know why.


[Edited on 23-12-2025 by semiconductive]

semiconductive - 22-12-2025 at 18:10

If there is a company who could answer this for a tutorial fee, I wouldn't mind paying out for a definitely correct answer.

Here on the forums, the discussion is free. But it's not like I'm cheap.

My college education set me back a few pennies. My ex wife set me all the way back. What's a few more $$, but this time it better be worth it.

semiconductive - 24-12-2025 at 13:49

Control experiment#2: Sulfite.

Since I can't replace the sodium with lithium, yet, I'm just going to replicate the same amount of alkalai in the last experiment but use sodium meta-bisulfite.

Afterthought:
I should have used lithium citrate + citric acid, but I just citric acid + lithium carbonate. So, I generate a little more water in this experiment which is bad.

Initial amounts:
3 [cc] ethanol 1 cc extra to allow more to evaporate
2 [cc] ethyl-citrate ( inert dilluting, and allows better visibility in coloidal suspensions. )
2 [cc] Kerosene To keep moisture and most air out of experiment.

Citric acid: 108 [mg]
Na-Met-bisulfite: 323 [mg]
LiCo3 125 [mg]

I was shooting for 0.0017 mole of sulfite, which gives 20-30 ethanol molecules for every sulfite. I might have messed up and gotten .0034 moles of sulfite.

Ground powder all together in a test tube using a glass stir rod.
MIxed in liquids.
Initial state was colloidal with clear kerosene floating on top, but it completely settled out in about a minute.

vlcsnap-2025-12-23-11h58m04s361.png - 350kB

Began heating between 70 and 85 [°C]. I turned down heat to stop boiling and paid close attention to any smells. There were no sulfur smells. It just smells slightly like ethyl alcohol.

vlcsnap-2025-12-23-12h08m17s459.png - 397kB

After running for about four hours, grey colored colloids began to circulate in the fluid. Black oxides stuck to test tube wall from the washer (steel anode) stack. Then color began to lighten, with some yellowish looking color forming on steel.

All brief color changes went away, and the solution converted to a very dark olive green after 8 hours. Green colloidal material settles out easily onto electrodes and tends to stick. WIth time it gets finer, and looser (less sticky).

vlcsnap-2025-12-23-19h09m50s494.png - 327kB

I will let this run a few days over Christmas and we'll see what happens.

semiconductive - 29-12-2025 at 12:38

I added 1/2 cc of ethanol to bring the polar liquid volume back up to about 4cc. Because of variability of liquid volumes after miscing, I am not sure how much ethanol is actually in solution any-more. I'm just attempting to maintain a constant volume and compare conductivity trends.

The color has remained a dark olive green. Over the last two days I have had small amounts of black material build up on the graphite electrode (-). I've cleaned it off. The majority of white and olive-green salt falls to the bottom of the test tube.

Each time I add ethanol, the temperature at which the bottom of the test tube forms boiling bubbles has dropped. Less material stays settled at the bottom of the test tube. I was able to run the test tube at 110 [°C] two days ago without boiling, but today I'm down to 80 [°C]. Pure ethanol according to literature boils at 78.3 [°C].

On the other hand, the conductivity of the solution has slowly risen in spite of the boiling point lowering.

I am running AC current, around 10 [mA] with a D.C. bias of maybe 1 or 2 milli amps.

This should cause fast and reversible chemical reactions to remain in solution, while more permanent or slower reactions will become electrode coatings (and I will clean them off).

Over time, I am hoping this will maximize conductivity of the solution by eliminating more stable oxides and traces of water. ( But I could be wrong. )

vlcsnap-2025-12-29-14h16m20s757.png - 238kB

Note: Rechecked it an hour after adding 1/2 [cm³] denatured ethanol.
Result, the bottom of the test tube is still olive green colored with white+green precipitate. But the top 3 [cm] of the tube is now browner shade.

I assume this is because introducing new moisture comes with introduction of alcohol.

For scale: I lost an eyelash into the test tube while scraping off the graphite electrode (- biased electrode ). The electrode itself is 0.9 [mm] wide. The view is highly magnified, with the glass test-tube oriented at 45 [degrees] to gravity. The top of the picture is level with the ground. Bubbles tend to leave the graphite electrode at a 45 degree angle, this way, and I'm able to tell how much gas comes from each part of the electrode.

In general, this is where hydrogen gas likes to form.

I should be breaking down water and ethanol at both the anode and cathode.



[Edited on 29-12-2025 by semiconductive]

semiconductive - 31-12-2025 at 13:02

For all intents and purposes, the solution has now become black uniformly black (6+ days of operation at 80+ [°C]. AKA: There's no point in posting more pictures. )

I will continue A.C. current for another week, and see if the solution clarifies as I scrape off residues from electrode.


semiconductive - 31-12-2025 at 23:30

Ficticious/non-isolatable chemicals problems:

In Alexander L. Shimkevich's paper on the electrochemical view of the "Band Gap" of liquid water, there are apparently differences in how Russian chemists theorize about dissolution processes vs. American and English authors.

In U.S. literature,
Dissolved hydrogen and oxygen are treated as a diatomic molecule that stays intact in water. But, Alexander is talking about neutral hydrated and dis-associated gasses in water.

Eg: Hypothetical O·H₂O = H₂O₂ molecule units with no charge. And, H·H₂O = H₃O molecules with no charge.

The former empirical equation is known to exist as hydrogen peroxide -- but the latter is not known.

Apparently, Russians have computed constants for these 'hypothetical' meta-stable H₃O states in a handbook of physical constants.

I am not completely able to understand Alexander's argument; but it's fairly clear that the energy gap he computes for H₃O⁺ vs. OH⁻ energy levels is 1.75 [eV] on page 245.

That is actually consistent with the maximum value that I computed using a curve fit a few posts ago. I computed 1.85 [eV] as the maximum the energy gap could be at absolute zero (Kelvin).

This energy gap is larger than a sum of two half reactions, pp. 246, (11) and (13) which are essentially standard hydrogen-oxygen fuel cell reactions, yielding 1.228 [V] at 25 [°C] and 1 [bar].

But, the equilibrium constants Alexander turned into voltages are very curious to me:

[ H₃O⁺ ]/[ H₃O ] = e^(( εH₂O - εF(2) )/(kB·T))
+0.219 [eV]

[ OH ]/[ OH⁻ ] = e^(( εH₂O - εF(3) )/(kB·T))
-0.302 [eV]

For, these imply that free hydrated and dissolved oxygen and hydrogen must be available in the liquid, and that they have a rather large total effect 0.521 [eV] on the difference between cell voltage and water's intrinsic energy gap.

Generally in reputable auto-ionization experiments, distilled water will have been triple boiled and is (therefore) automatically de-gassed. There really ought not be any dissolved gasses in the data I've been curve fitting, unless it's introduced by hydrogen electrode apparatus. (No oxygen would be introduced, though).

When I used the half cell reaction for hydrogen gas and hydroxide ions and tried to force the energy gap to be 0.828 [eV], I was thinking that a SHE + a graphite electrode in pure water explained the entire situation. But now I realize, my earlier attempts were over-simplified.

If I'm understanding what I'm reading in various chemistry texts:
The Nernst equation, and the Gibbs equation, are based on the ideal gas law relationships.

But, liquids deviate from these laws based on vander-Wall's forces.
Which means the temperature dependency of the Gibbs equation can be influenced by chemical properties of water.



[Edited on 2-1-2026 by semiconductive]

semiconductive - 2-1-2026 at 21:01


--- Semiconductor fitting of water auto-ionization, what worked, what didn't ---

Reviewing the NIST™ data and comments, it's fairly clear that the density at both 0 [°C] and 100[°C] are very difficult to measure accurately. The experimental data is only guaranteed valid over the range of 5 to 40 [°C].

Note: Online sources often use the NIST density polynomial at 0 [°C] as if it
were an accurate value, which I have discovered is a big mistake.

When comparing Molal pH+pOH values, vs. Molar pH+pOH values, the conversion does not significantly change my plot except at zero degrees celsius. 0.1/T=0.366... on the X-axis. I have a lot of reasons to believe this data point is wrong after running many fits. This point is a major outlier!

-----

When I plot water ionization curves with 1/T [Kelvin] as the x axis, and pH+pOH as the y axis, the slope of the curve represents the physical property of 'energy required to make ions'.

By inspection of the Arrhenius plots that I've already done, it is evident that water auto-ionization slope is always less steep on the left side of the plot (where the liquid is hotter) and more slanted on the right side of the plot ( Where liquid is colder. )

Therefore, the energy ionization gap of H₃O⁺ to OH⁻, decreases with an increase in temperature. This physical behavior is identical to crystal based semiconductor behavior. The energy gap is largest in coldest materials.

Notice: in liquids the colors of complexes are often computed using crystal field theory, even though water isn't really a crystal with fixed length bonds.

But: After some experimentation, I realized that Semiconductor physics are derived assuming a fixed sized crystal lattice and thermal expansion is typically handled by Varshni's correction.

Therefore, to get the same model accuracy with water, it's necessary to convert the ion density of water into a Molal measurement basis rather than a Molar basis.

I used least squares fitting of all data, including the erroneous/outlier data point at 0[°C], on a Molal basis, to get the following equation:

pH+pOH (Molal) ≈ 4674.12/T -10.5102 +0.651441·ln(T) + 7.4969·10⁻³·T²/( T+1.36780)/log₁₀(e)

This has a maximum energy gap of 4674.12 / 2519.88 ≈ 1.855 [eV]
( Which is reasonably close to another paper's calculation of it's value -- except that it's wrong by a factor of 2. ).

Notice the fit equation's intercept value is large at -10.5102; and this value is a correction of sorts, because if the model was exact the intercept should be nearly zero.

On the other hand, The Varshini correction is much less than '1'. In my experience, the intercept correction is less important than the Varshni correction because the intercept is affected by effective masses.

I got a much worse fit when fitting only 0 to 40 [°C] data, in an earlier post.
However, if I run data points from 5 degrees to 40, and purposely exclude 0 [°C]; I get an extreme improvement that is even better fit than the full data set.

pH+pOH (Molal) ≈ 4615.39/T -10.0901 +0.651441·ln(T) + 7.16134·10⁻³·T²/( T+1.06895)/log₁₀(e)

Max energy gap: 4615.39/2519.88 = 1.831 [eV] # Fit of 5 to 40 [°C]

I can even re-check using data points that are in the 'too high' range, and they produce about the same result:

max energy gap: 4522.85/2519.88 = 1.794 [eV] # Fit of 40 to 75 [°C]

This indicates that the ionization data point at 0 [°C] is an outlier.

---- Correcting for effective mass without having a plot of the QM E-k rleationship ----

Intrinsic ion count in a semiconductor is usually written as:
n_i² = Const · T³ · ( m_e · m_h )^(3/2) · e^(( Ev-Ec)/(k·T))

Where (in silicon) m_e, m_h are the relative mass multipliers for electrons and holes. Ev-Ec is the energy required to ionize an electron from valence orbital into conduction bands.

I have been fitting a square rooted version of this equation, where the effective masses are both considered to be unity.

n_i = Const ·T^(3/2) · (1·1)^(3/4) ·e^(E/(2·kB·T))

Which is a mistake on my part.

I forgot that n_i is the value of one kind of carrier, only, eg: pH or pOH by itself.
The carrier concentration product pH + pOH = -log₁₀( n_i² ).

When I do a best curve fit letting the temperature power be variable, I get a best fit of Kw data with an exponent of 3.02, not an exponent of 3/2. This verifies that I made a mistake.

Kw is equivalent to n_i², not to n_i.
The correct fit for the equation 5 to 40 [°C] is:

pH+pOH ≈ 4327.11/T -1.02766 - 3·ln(T)/ln(10) + 5.9858·10⁻³·T²/( T+1.022)/log₁₀(e)

Going from 5 to 95 [°C], gives almost the same:

pH+pOH ≈ 4293.81/T -1.24649 - 3·ln(T)/ln(10) + 8.44233·10⁻³·T²/( T+1.017)/log₁₀(e)

But, this also means the energy gap scaling for the earlier part of the post is wrong by a factor of 2.

The remaining (tiny) errors in my plot are due to two issues, that semiconductor derivations of density of states (DOS) don't use the exact Planck distribution because it has no analytical solution, Rather the derivations I've been linking substitute in a Boltzmann approximation.

The Boltzmann gas approximation is very accurate as long as the ionization energy gap is large compared to the the thermal voltage of the ions: eg: 6·kB·T > Energy of ionization. Thankfully I know that 6·0.024 [eV] = 0.144 [eV] at 0 [°C].

Since all ionization energies that I've computed are *easily* bigger than 0.144 [eV] , even at freezing temperatures, I know the standard semiconductor approximation isn't introducing any significant error to the properties of water.

The second source of errors is that Density of States formulas were derived assuming that he speed of light, 'c', has a fixed value. But: liquids and solids have lower values for 'c' than empty space does and (worse) they change with temperature.

In a true semiconductor setting, I would work out an E-k (energy momentum diagram), based on atomic orbitals and use that to compute an 'effective' mass for the ions that accounts for the change in the speed of light.

But, that approach is impractical here and the error that is being corrected is very small (typically much less than 1%).

Since, I already know that both classical and relativistic energy-momentum diagrams can generally be described qualitatively as a hyperbola, I'm going to attempt to fit a simulated E-k diagram to water auto-ionization data.

From the standard semiconductor equation, it ought to be obvious that the DOS mass equivalent is a geometric mean between two ion masses.

n_i² = Const · T³ · ( m_e · m_h )^(3/2) · e^(( Ev-Ec)/(k·T))

Therefore, I can replace the geometric mean with a single effective mass that is a function only of temperature:

( m_e(T) · m_h(T) )^(3/2) → ( m_DOS(T) )^3

The effect on pH + pOH is an error proportional to: 3· log₁₀( m_DOS(T) )
Therefore: m_DOS ≈ 10^( pKw_error/3 )

I am only fitting Kw data from 5 to 40 [°C], but even so the extrapolation of the curve is excellent even to boiling.

pngsnap.png - 14kB

The quality of the data from 40 to 100 [°C] is unknown. The reason is that the international temperature standard ITS-90, document, indicates that two different experiments might be involved in making these measurements. eg: The water samples and equipment making measurements are not necessarily the same in these two temperature regions.

On the plot, the break point is x=0.0032. There's obvious trend-changes around this temperature, and that could easily be caused by interpolation errors from different data sets or slightly different compositions of water.

The relative permittivity equation that I found earlier when researching "Kell", is clearly an optical permittivity. Trying to re-find the equation using google searches doesn't work. I wonder if I've been given tampered with documents.... sigh.

When I compare it with other published data points, I don't get good agreement.

pngsnap.png - 13kB

But, I need to estimate the speed of light for the energy being applied to the ion in order to do a theoretical effective mass correction.

If I divide the index of refraction's reported experimental data by the cube root of density of water that NIST publishes for ITS90, I get straight lines. But you can tell which lines are theoretical because they actually curve slightly.

For all practical purposes, light with energies of 0.5 [eV] to 6 [eV] have a constant delay per molecule of water encountered! This will simplify the work I need to do to compute effective mass changes over temperature for my semiconductor model.




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[Edited on 4-1-2026 by semiconductive]

semiconductive - 5-1-2026 at 15:36

Hmm. I need to be able to re-do this derivation with alcohols, and I need it to be simple.

The NIST models are too complicated for hand computing without an online calculator. I'll work out a restricted standard conditions case, here.

I pretty much just need to know the refraction properties of liquid (water, alcohol) at 1 [bar] pressure for my own experiments.

I also need to verify what the density of water is over the full temperature range, and not just 5 degrees to 40 degrees Celsius.

I can get high quality data for index of refraction in water from multiple sites.
Therefore, I'll test out a theory out about how the index of refraction changes vs. temperature.

Using just the index of refraction data (mostly appendix info) from:

https://www.researchgate.net/publication/252888306_Water_ref...

And from NIST:
https://srd.nist.gov/jpcrdreprint/1.555859.pdf

I hypothesize:
Index of refraction is nearly constant after dividing out optical path length changes that are caused by water density. Eg: the delay in light travel time is directly proportional to the number of molecules that are encountered independently of empty space traveled through. Therefore, I suspect that the primary systematic error in measuring refraction experiments ( AFTER volume phenomena are removed/compensated for) will be inaccuracies in temperature.

That means, if I normalize the index of refraction for a color passing through water , by dividing by the index of refraction of a specific color, eg: 589.32 [nm] light at the same temperature, the implicit volume changes get canceled out because both refractions values have the same volume/density of water.

But: Slight errors in temperature or in-homogeneity of water mixing, will result in linear slope errors. IF this is the case, then dividing the index of refraction of light by another refraction index of light of a different color will leave gently sloped lines that indicate temperature mismatches/errors and other linear energy change errors.

*These errors should be easy to spot in replications of the same experiment by different authors.*

pngsnap.png - 14kB


YES! The plot looks very linear, as expected!

The straight line least squares fits are:

1.04533 -18.8847·10⁻⁶ · (t-25) @ 226.5 [nm]
1.01131 -3.69776·10⁻⁶ · (t-25) @ 361.05 [nm]
1.00771-2.02039·10⁻⁶ · (t-25) @ 404.41 [nm]
1.00751-8.29188·10⁻⁶ · (t-25) @ 404.66 [nm]
1.00039+1.69418·10⁻⁶ · (t-25) @ 589 [nm]
1.0 @ 589.32 [nm]
0.997912-2.23743·10⁻⁶ · (t-25) @ 706.52 [nm]
0.994315+6.72551·10⁻⁶ · (t-25) @ 1013.98 [nm]
0.957622+37.0681·10⁻⁶ · (t-25) @ 2325.42 [nm]

The standard reference for 589.32 [nm] light at 25[°C] and 1[bar] is n=1.3325
Therefore, my linear fits predict at 25[°C], 1 [bar], vs. "simple approximation", vs. NIST's recommended fit.

n= 1.3929 @ 226.5 [nm] vs. 1.3230 vs. 1.3925
n= 1.3476 @ 361.05 [nm] vs. 1.3441 vs. 1.3474
n= 1.3428 @ 404.41 [nm] vs. 1.3415 vs. 1.3426
n= 1.3425 @ 404.66 [nm] vs. 1.3415 vs. 1.3426
n= 1.3330 @ 589 [nm] vs. 1.3322 vs. 1.3328
n= 1.3297 @ 706.52 [nm] vs. 1.3288 vs. 1.3299
n= 1.3249 @ 1013.98 [nm] vs. 1.3242 vs. 1.3248
n= 1.2760 @ 2325.42 [nm] vs. 1.3204 vs. 1.2758

The simple approximation is close everywhere except the two extreme wavelengths.

But: The NIST equation, #7, although accurate for my data set -- doesn't even reproduce the first wavelength between the values listed in the NIST tables for 20 to 30 degrees of wavelength .36105 micron in it's own document. The author has tables that aren't computed by the formula he publishes ?

To represent standard laboratory conditions at 25[°C], (1 bar pressure) I set:
T=298.15; T₀=273.15
ρ=997; ρ₀=1000
λ=361.05; λ₀=589.0

Using Table 4 in the document for full range optimized variables, I compute with equation (7):
n = 1.3474
Using the different coefficients in the Appendix, I still compute,
n = 1.3474

This value is not between n=1.39336 at [20°C] and 1.39208 at [30°C] on page 704 in the 0.1 Mpa pressure column. (1 bar).

I really ought to use experimental data, first... But now comes another problem.... index of refraction values at standard temperature and pressure are contradicted by different published sources:

Kudos, Philiplaven, for pointing this out!
http://www.philiplaven.com/p20.html
Note: My estimated values are at close to IAPWS & Lynch & Livingston's.

Alternately, I thought I could use Wikipedia's tables:
https://en.wikipedia.org/wiki/Optical_properties_of_water_an...
But my values substantially disagree. (Jan 2026)

What a nightmare! This is basic science.
Curve fits are fickle....

pngsnap.png - 28kB

NIST = Journal of Physical and Chemical Reference Data 19, 677 (1990); https://doi.org/10.1063/1.555859 19, 677

--------- Side Note:
Optical absorption of light in water hits a maximum at around 2.7 microns. This light's energy ought to roughly correspond to the energy gap of water OH⁻, H₃O⁺ or an integer fraction because this is the primary energy being used to cause ionization of water.

Energy of 2.7 micron light is 0.459 [eV].

This value corresponds closely to the article that BNull linked, regarding frozen water diodes made of dilute acid and bases.


[Edited on 6-1-2026 by semiconductive]