Sciencemadness Discussion Board

Easy sulfite ion in a pinch...

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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 75 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 49 times

Attachment: S. Yefimov - Band gap of water.pdf (292kB)
This file has been downloaded 51 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.




pngsnap.png - 12kB

[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]

semiconductive - 9-1-2026 at 09:57

I've asked "'Elena Genina", if perhaps there is a misprint in her article for a "simple fit" of coefficients because her index of refraction doesn't match NIST/IAPWS or their own data near 200 [nm]. Until I get an answer, I'll just set that equation aside.

I can regress my normalized index of refraction values to any temperature; at 20 degrees Celsius, I note the original formulas Elena researched become valid.

Experimenter "P.O. Rol's" fit ignores the infra red, but still agrees well with what NIST/IAPWS predict for the same temperature in the visible to UV. His fit does not have a defect in 'n' near 200 [nm].

pngsnap.png - 31kB

Note: NIST and IAWPS can't be told apart in the plot, they follow the same line.

Out of the nine experiments which I linearly regressed, there are four that are nearly duplicated pairs: 404.41 [nm] vs 404.66 [nm] and 589 [nm] vs. 589.32 [nm]

These four experiments are not near the troublesome resonance wavelengths of a "Sellmier" model as used by NIST/IAPWS. Therefore, the slopes of these four equations aren't affected much by small changes in color even over large temperatures. The same is not true of the near UV or medium wave IR light.

I'll work with these four experiments, as they are simpler, to get an idea of how the index of refraction behaves.

Note: When I take a linear regression to 20 degrees, the data points I get agree with P.O.Rol's fit even better than with NIST/IAPWS's prediction at 20 degrees.

There are three theories that index of refraction is predicted by, typically,
Kramers-Kronig, Lorentz-Lorentz, or Gladstone-Dale. Each has surprising refinements in terms of volume.

I find this curious, because:

If I take the cube root of density of water, the number I get is equivalent to the length of one side of a water-cube which light is passing linearly through. Since the number of molecules in the volume is fixed, changes in the length of the cube correspond to changes in empty space between molecules where light is free to travel at it's maximum speed.

By definition: The index of refraction ,n, corresponds to the total amount of time taken for light to travel a fixed linear distance.

Therefore, if each molecule (at a given temperature) were to delay the passage of light by some fixed/average amount of time, each; then I expect the index of refraction, n, to be a simple weighted average in proportion to the number of molecules encountered per unit length.

But, none of the models of refractive index are based on the cube root of density.

The simplest relationship appears to be in the article from March 5,1863 -- "Researches on refraction, dispersion, and the sensitiveness of Liquids"
J.H. Gladstone PhD, and rev. T.P. Dale. ( Royal Society publishing org. )

Taking the index of refraction, subtracting one, and then multiplying by the volume ( AKA divide by density ), generally produces something very close to constant.

The calculations done in the original article are only aware of the Cauchy, Hamilton formula:
n = A + B/λ² + C/λ⁴ + D/...

But from the linear nature of my regression fits, I'm pretty sure that the color of light that will have the least deviation will be the one where the curvature of the dispersion relation momentarily goes to zero -- between the UV and the IR resonance.

I can run a numerical test on the NIST/IAPWS curve to approximate the inflection point wavelength vs. temperature. From the previous plot, I estimate that the inflection point is at about 850 [nm] at 20 [°C].

The Gladstone/Dale relationship works with alcohol and acetone, as well.


[Edited on 10-1-2026 by semiconductive]

semiconductive - 12-1-2026 at 15:29

My estimate, by eyeing the plot, was off. The color of light that is most linear in terms of volume/temperature ought to be ≈1003 [nm].

Studying the Lorentz-Lorentz equation used by NIST and IAPWS, I notice that the density variable is not multiplied directly by the temperature variable anywhere. Temperature and density are handled independently in the equations.

This independent handling means I can plug a wrong (but constant) density value into the equations and I'll get the same inflection point as if I plugged in correct density values. eg: I will only change the 'n' value by a fixed amount, and therefore the error won't affect computation of the inflection point wavelength where index of refraction changes from positive to negative curvature.

I get the following answers regardless of whether I plug in a constant density, or ITS-90 polynomial.

[°C] n-NIST n-R9-97
10.0 998 997 [nm]
15.0 999 998 [nm]
20.0 999 998 [nm]
25.0 1000 999 [nm]
30.0 1001 1000 [nm]
35.0 1001 1000 [nm]
40.0 1002 1001 [nm]
45.0 1002 1001 [nm]
50.0 1003 1002 [nm]
55.0 1004 1003 [nm]
60.0 1004 1003 [nm]
65.0 1005 1004 [nm]
70.0 1005 1005 [nm]
75.0 1006 1005 [nm]
80.0 1007 1006 [nm]
85.0 1007 1006 [nm]
90.0 1008 1007 [nm]

Doing a little more research, the Gladstone-Dale relationship is a dilute solvent limit to the Lorentz-Lorentz equation.

(open access pdf).
https://pubs.acs.org/doi/10.1021/acs.jpcb.5b05433

Therefore, I can use either equation for auto-ionization calculations. They will be equally accurate. Gladstone-Dale being simpler, is preferable.

The index of refraction is usually reported to 6 digits; but the temperature's accuracy in these experiments isn't likely that accurate. I expect ±.1 [°C] as an excellent experiment but not six digit temperature accuracy.

To measure density based on refraction:
The closest wavelength I have to the ideal ≈1003 [nm] is 1013.98 [nm].
My linear regression is indistinguishable from NIST and IAPWS values at 20 [°C], so I think this wavelength is sufficient for my purposes.

I ought to be able to use density values from ITS-90 for 10 to 40 degrees [°C] to compute idealized data slope for a Gladstone-Dale fit to 1013.98 [nm] data points.
I can then compute a best Gladstone-Dale fit for experimental data which has the same ideal slope. The differences between the fit and the actual data will allow me to estimate temperature errors.

Hopefully, I'll get a nearly Gaussian profile of errors and can then trust the experiment to infer what water density is over all liquid temperatures and not just 5-40 [°C].

I'll also compute the fresh water fit of Quan and Fry, 1995, divided by 1.003 because the equation doesn't depend on volume calculations from ITS-90.

However, the fit range for oceanography generally doesn't go above 30 [°C].
https://www.oceanopticsbook.info/view/optical-constituents-o...

T [°C] G-D(NIST) G-D(IAPWS) G-D ( Quan-Fry )
10 0.32600 0.32600 0.32629
20 0.32582 0.32582 0.32613
30 0.32566 0.32565 0.32591
40 0.32550 0.32550

Slope of Gladstone-Dale "sensitive energy", for Quan-Fry, NIST, and IAPWS
-18.700 · 10⁻⁶ -16.600 · 10⁻⁶ -16.700· 10⁻⁶

Average linear Gladstone-Dale intercept @ 0 [°C] = .32627
Average linear slope = -17.333 · 10⁻⁶

With this averaged Gladstone-Dale fit, I can now estimate the density of water that was required to get the experimental data at 1013.98 [nm]:

pngsnap.png - 15kB

And, it's pretty obvious that the ITS-90 standard/Kell equation is not very good outside of the 5 to 40 degree Celsius range. At 90 [°C], the Kell equation is 0.9620 [ g/cm³ ]. My Gladstone-Dale redacted water density is of about 0.9656 [ g/cm³ ] at 90 [°C].

For the Kell equation to be correct, the experiment's thermometer would be in error by a full 4.2 [°C]. That's extremely unlikely!

eg: AI's telling us that KELL is the most accurate fit, is misleading.
Kell is only the best fit over part of water's liquid temperature range.

IAPWS, for industrial standard -- shows 0.96531 [ g/cm³ ] -- at 90 [°C]. Therefore, even the supposedly less accurate "industrial standard" is closer to scientific experiments than Kell at 90 [°C].

Well, I learned something. ITS-90/Kell water density polynomials are not sufficient for general chemistry. Ouch.


[Edited on 13-1-2026 by semiconductive]

semiconductive - 14-1-2026 at 21:04

I'm not really making progress; The plots I am getting from IAPWS water density polynomial make me suspect that refractometers may have severe accuracy problems measuring liquids above 50 [°C].

I am not able to find commercial refractometers available for less than $6000 U.S. that can measure six digits of precision. It's out of my budget range. But, on closer inspection -- only the Milles Griot "M4" model boasts the ability to make measurements over a full 0 to 100 [°C] range, and not at 6 digits precision.

That means the equipment to make measurements of water refraction over wide temperature ranges are not common, industrially.

Looking around for how different people attempt the feat, I see one creative solution by Kendir & Yaltkaya where they make a custom device to measure index of refraction over temperature using fiber optics. K&Y only claim 5 digit index of refraction precision.

When I plot the Gladstone-Dale relationship for K&Y's experiment; I've chosen to use the standard specific density polynomial from IAPWS-97 against their data fit.
See result below: I get quite a wavy line. There are errors of larger than 5 degrees Celsius in the non-linearity.

Their result is also qualitatively different than when I plot Bashktov & Genina's raw data against IAPWS-97 density. I see reasonably linear Dale and Gladstone data over 0 to 40 [°C], but then the plot gently curves upward after 50 [°C].

This suggests that the refractometers are not all operating as expected.
Note:
Using linear regression, I estimate maximum B&G temperature errors must be two degrees Celsius in order to explain the remaining non-linearity if the refractometer was operating properly.

That's not an impossible amount of error for a thermocouple...
But, I have trouble believing that all experiments have at least that amount of temperature error.

I am not going to be able to extend these results to alcohol, if I can't even get them to work with water which is extremely well documented. :(

I'm not sure how to proceed as I have no experience trouble-shooting refractometer experiments.

pngsnap.png - 44kB

I used the following specific density values in the plot, because the IAPWS Python library yields these specific values for water density:

[°C] @ 100 [ kPa ] @101.325 [ kPa ]
10 0.9997009 0.9997015
20 0.9982055 0.9982061
30 0.9956515 0.9956521
40 0.9922237 0.9922243
50 0.9880469 0.9880475
60 0.9832100 0.9832106
70 0.9777787 0.9777793
80 0.9718023 0.9718029
90 0.9653181 0.9653187

Things I have learned:

The IAPWS difference in volume between 1 bar and 101.325 [kPa] do not affect the plotted errors significantly. The difference in volume between de-gassed water and fully saturated water do not affect the plots significantly.

Replotting the error of the IAPWS theory converted to a Gladstone-Dale product gives about half the error. It notably curves gently upward above 50 [°C] replicating the qualtitative aspect of the raw B&G data-points, correctly.

pngsnap.png - 50kB

The extremely large jump in values from 0 to 10 [°C], shows that freezing point is still an outlier, even when using purely theoretical values from IAPWS to compare against experiment.

I (presently) imagine three possibilities that might explain the non-linearity of the plots: 1) Abbe refractometers do not measure the speed of light in liquid accurately at higher temperatures due to some systematic source of error. 2) There are inhomegenities in water where the liquid is a mixture of micro-boiled pockets, liquid pockets, and frozen or paired molecules with different refraction indexes. 3) There is an unknown amount of impurity (salt?) in the water of some experiments and not others.

In a previous post:
When I divided the index of refraction of one wavelength of light by another to 'normalize' the index of refraction, I did this to cancel out volumetric non-linearities between colors of light. But, it's also true that mixed phases of liquids could also tend to cancel out since the mechanisms are similar.

I imagine mixed phases in fixed proportions might have the same effect as changing the volume by a definite amount; if so, then I can imagine a 'virtual volume' exists for distilled water at every temperature. This virtual volume can be different (slightly) from it's physical volume -- but it will cause Gladstone-Dale plots to become linear over all temperatures.

My hope comes from the previous post I did:
The extreme linearity of a wavelength *normalized* G-D plot (except at extreme colors near water resonances), makes me think that regardless of what systematic error there are in abbey inferometry, those errors might be cancellable empirically.

eg: the most difficult to remove errors between refractometers is probably caused by manufacturer errors in the reference scale or the zero angle being off slightly. These could only be totally detected when the exact same experiment is repeated using different devices with the same liquid sample. I don't have that luxury.

But, I can make an unjustified assumption which may be sufficient to overcome the problem. I can assume that any linear offset and rescaling of experimental refraction data that improves the over-all fit is removing more error than it is introducing.


[Edited on 16-1-2026 by semiconductive]

semiconductive - 16-1-2026 at 20:07

A closer look at the nature of the errors:

In 1990, a change was made in the practical value of temperatures reported in literature. In the 0 to 100 [°C] temperatures reported before 1990 supposedly need to be divided by 1.00024 to match data published during or after 1990.

The only exception is temperatures reported for oceanography, which supposedly continues to use the older temperature scale.

For refraction of light purposes:
I know the speed of light was defined as a constant value back in 1983.

I've only been plotting data from articles after 1983; Which means the reference speed of light ought not have changed between physics journal articles.

"Journal of Chemical Reference Data, Vol 14, No.4, 1985"
https://srd.nist.gov/JPCRD/jpcrd282.pdf

Bashkatov and Genina (2003) were quoting index of refraction data identical to:

"Journal of Chemical Reference Data, Vol 19, No.3, 1990"
https://srd.nist.gov/jpcrdreprint/1.555859.pdf

Unfortunately, there is a misprint in the wavelength title of the 1990 journal on page 704. JPCRD labeled tabulated data as being 361.05 [nm] when it's clearly 226.50 [nm] data. I noticed that the formula in this article does NOT produce the values near p. 704's values when 361.05 [nm] is plugged into the formula.

Bashkatov and Genina correct the erroneous label in their "simple model" article, but for reasons I have not understood -- their fit still doesn't produce the correct 226.5 [nm] data when 226.5 is plugged into their fit equation. Whether I made a mistake, or they fit the wrong data, I don't know. But, I can't use their fit.

With that in mind, here is a Gladstone & Dale plot of undisputed data from both reference journal articles in 1985 and 1990:

I have converted the 1985 Celsius temperatures on the plot to agree with 1990 standards by dividing them by 1.00024. I have also corrected densities at the offset temperatures of 1985 data by using the IAPWS density at a temperture divided by 1.00024.

Even with these two corrections, there is a clear difference in the plotted quality of 1985 tabulated data vs. 1990 (and later) tabulated data.

pngsnap.png - 14kB

Note: The grey lines are the theoretical G-D values using IAPWS formulas for all wavelengths and temperatures and volumes. These lines tend to curve upward, slightly after 45 degrees Celsius compared to a straight line fit. The theoretical formulas are within 1/3 of a degree error compared to the 1990's table values.

The same is not true of the 1985 data, even after temperature and volume correction. The index of refraction is uniformly low by an amount that would require an unacceptable temperature error correction on the order of ~10 degrees Celsius everywhere.

Additionally, the slope of the line through the data points curves *down* after 45 [°C] rather than up.

I have tried every conceivable change between STP (1 ATM) and (1 BAR), in addition to reversing the ITS90 change to ITS68. None of these changes affects the position of the 1985 data or its' slopes significantly enough to make a harmonized plot. The difference in data isn't due to a simple mistake in conditions being recorded wrong.

Something is systematically different (and un-documented) about the standards used for the index of refraction values published in 1985 compared to 1990 and after. Correcting temperature, pressure, and volume for the ITPS68 to ITS90 standards change is not enough to make the tabulated data agree from these different time periods.

Here's a plot where I multiply the index of refraction of the 1985 data by the average calibration scaling error: n_1990 ≈ n_1985 · 1.00046

pngsnap.png - 14kB

Notice the average index correction .00046 is roughly double the temperature change correction 0.00024 required to convert between ITPS68 and ITS90.

With exception of the yellow sodium line (589.32 [nm]), which has noticably smaller error than all other 1985 data line plots; It's obvious that the tail non-linearity trends are opposite in 1985 vs. 1990 plots.

I have changed the 1985 index of refraction scale, and the only other calibration possibility is that the investigators had a calibration error that introduced a constant offset. But no constant offset of index for each individual experiment is capable of removing the remaining errors.

All 1990's data fits and theoretical equations are self consistent within less than 1 degree Celsius error.

1985's data shows consistent errors and trends in all experiments that are far larger than can be explained by a tiny temperature scale change in 1990.

I know that if I were to use G-D plots to solve for water volumes for 1985 and 1990 tabulated data by requiring straight line line fits, what I would end up with are two distinct and mildly dis-agreeing density curves. One that works for 1985 data, and the other which works for 1990 data.

I can imagine one way that such an error could have gotten published historically.

I suspect the tabulated data in the physics journals are statistically smoothed data in the lower digits. The tables are not truly raw measured data points. Whatever systematic error and problems Abbe refractometers introduce, these problems have not changed between 1985 and 1990. Therefore: The same kind of error ought to show up in the data for both time periods, and doesn't.

I think this means the 1985 authors used a density curve for water that slightly disagrees with the one used in 1990 when statistically smoothing their data.

eg: The density curve for ITS-90, based on Kell's data from an earlier time period, is only valid in the 5-40 [°C] region. Therefore, we (as international scientists) don't have any evidence of what volumes were used by NIST in 1985 for the remaining 40 to 100 [°C] region.



[Edited on 17-1-2026 by semiconductive]

semiconductive - 18-1-2026 at 22:09

The last test tube turned completely black over time.
I'm curious if it's the citric acid that makes a black chelate in alcohol.

There's a few percent water in normal ethanol. I think that's what is making the rust in the first place. So, I'm going to see if I can get it to precipitate out with glucose. Glucose supposedly can sequester and/or chelate iron oxides. Glucose isn't very soluble in ethanol to begin with, so I think maybe saturating the solution with sodium metabisulfite and glucose will precipitate out the orange/black oxides while water is being broken down.

1 CC -- ethanol (Fuel grade).
3 CC -- tri-ethyl-citrate ( An inert diluent. )
1 CC -- kerosene ( To float on top and keep air/moisture out. )

The anode from the last experiment had a little black oxide on it, which I didn't bother to clean off. I added 19 fresh steel washers in a spaced stack to make a lot of surface area.

The bottom of the test tube is held at 80 [°C].
Glucose ~50 [mg]. ( Tiny dusting. )
Sodium Meta-Bisulfite -- enough to where it stops dissolving and leaves 1/3 cc volume of crystals in bottom of test tube.

Glucose initially out-gassed while dissolving in 1CC ethanol. Colloidal glucose crystals can be seen dispersed in the solution, not fully dissolved.

Rust dissolved off the anode and turned solution a faint orange, initially.

The current level is very low, 50 [μA] x 18 plates ≤ 1 [mA].
Hydrogen can be seen bubbling off the cathode, slowly, after 24 hours.

vlcsnap-2026-01-18-22h01m07s544.png - 281kB

The solution has gone from orange to yellow, and is clarifying slowly. Edit: initially, there is precipitation of brown stuff into the sodium meta-bisulfite, but it saturates after about 12 hours. The clarification might never complete. I would have to centerfuge it, to remove colloids, and add more glucose to definitely precipitate all oxides.

I'm hoping that the fact that it isn't turning black (yet), means that the dark color of the last experiment may have been caused by lithium citrate in the presence of moisture.

Glucose has an aldehyde bond on the end, but otherwise is alcohol like. I don't expect it to turn black. Hopefully, glucose will help retain ethanol in the tube at higher temperatures.

Edit: I wiped off the graphite electrode this morning after noticing that it was heavily bubbling hydrogen even though the current level hadn't changed since last night. After putting it back in , no bubbles formed immediately. Then slowly gas began building up and bubbling faster over a period of hours. There's obviously something (probably glucose) building up on the surface of the electrode that reacts to release gas.

Since I see no gas forming at the bottom anode of the 18 washer/plate stack, I decided to raise the temperature of the test tube until I could see one or two tiny bubbles form a minute at the bottom. The test tube bottom is now held at 105 [°C] and the kerosene on top is 30 [°C]. Extra heat at the bottom of the test tube is not significantly changing the gas production rate at the anode.

: Speculation :
This temperature is above the boiling point of ethanol at the bottom of the test tube, but since the top of the test tube is way below the boiling point of ethanol; I expect the washer stack to act like a distillation column and re-condense ethanol before it evaportates.

Edit: 6 Hours of 105 [°C], and I've lost roughly 0.1 [cc] of liquid. The test tube is making bumping sounds. I turned heat down to 101 [°C]. It's still bubbling from the bottom, but it's quiet now.

After adding back 0.1 [cc] of ethanol, the current jumped up to 200 [μA], 200x18 plates ≈ 3.6 [mA] chemical oxidation and reduction activity.

That's a reasonable level of current.

Edit: Had to lower temperature again, 95 [°C]. When I did so, glucose fell out of solution rapidly leaving colorless liquid with small orange floating specs. Allowed liquid to settle in a transfer pipette, and ejected brown/orange sludge into trash. Added fresh glucose and Sodium meta-bisulfite, shook, waited for it to settle and then decanted off the liquid. Now I have clear liquid with white floating glucose specs.

Current level was un-affected by cleaning, and even by sanding the washers lightly.
I added another 0.1 [ml] of ethanol, and current only rose 20 [μA] AKA 5%.

I then decided to experiment with other alcohols.
I added 3 drops of 1,3 propanol. Current level only rose another 20 [μA] over an hour then stopped rising.
I added 3 drops of glycerol, Current level rose only 10 [μA] over another hour and then stopped rising. Glycerol was not entirely dissolving and left a whitish layer on the anode. I mixed it in, and then noticed more bubbling on cathode with a little bit of shiny material depositing (probably glycerine). I waited another hour, no current change.
I added 3 drops of methanol. Current level went from 225 [μA] to 475 [μA], 211% increase. Gas bubbling out over doubled.
See picture:


vlcsnap-2026-01-19-18h33m04s407.png - 277kB

The yellow that you see is a stained Nylon standoff. The actual fluid, is not colored, but has small floating white specs in it. I probably need to get a micro-filter that can do a test tubes worth of liquid. Decanting isn't very effective.

The glucose worked far better than I expected. :)



[Edited on 20-1-2026 by semiconductive]

semiconductive - 20-1-2026 at 12:25

Re-run of experiment without glucose, in AC current mode, to get a better idea of the side reaction happening with 1,2,3-propanol.

There is a DC bias current that I measure, but note the major current is an un-measured symmetrical AC current. This combination of currents is to make reaction speeds more visible on the steel washers since the bottom side of each washer acts as an cathode more often/longer than it acts as an anode.

Note: My plastic droppers drip about 25 drops of water to reach 1 [cc].
It won't be exact for alcohol...

1 [cm³] 1,2,3-propanol 98% Duda-Energy™
3 [cm³] ethyl-citrate reagent grade
~100 [mg] Na-metabisulfite reagant grade

Initial current was a pathetic, 2 [μA]. The PP-ol was all sunk to the bottom, happily dissolving the metabisulfite.

I added 3 drops of methanol to see if I could get it to misc, better. The current went up to 200 [μA] just like the last experiment. But, then current began fading. A lot of gas was being generated.

I could have used ethanol, next, but I wanted something that would misc better with ethyl-citrate than with glycerol and act sort of like a soap.

I chose iso-amyl alcohol. 3 drops.
Current went back up to 175 [μA] but began dropping a bit slower than with methanol. So I added another 6 drops. It held on for a half hour before starting to drop back down.

So, I opted to put three more drops of methanol in, and the current went way up to 550 [μA]. After seeing it was stable, I polished the steel electrodes with sandpaper, mixed the solution well and added another 6 drops of methanol. Current rose to 800 [μA] and started climbing. (!! I went to bed !! )

vlcsnap-2026-01-20-12h10m59s755.png - 316kB

This morning the characteristic ?glycerol? side-reaction deposit showed up only on the final steel electrode. ( Most cathodic voltage )

Black oxide-like deposits built up uniformly on the bottom side of all other washers, with no significant glycerol side product. This suggests that either water or glycerine makes a black iron when being reduced/deposited.

Discussion:

The current level is pretty much back to what it was after adding iso-amyl alcohol.
Since this is an AC experiment, I expect ions ought to be building up in solution with any slower side-reactions remaining on the washers.

But, the current levels suggest that hardly any extra ions exists now.

I think this means that the methanol is being consumed and probably turned into a gas or an inert liquid.

pngsnap.png - 35kB

[Edited on 20-1-2026 by semiconductive]

bnull - 21-1-2026 at 12:08

Is there any strange smell, something vaguely like cookies?

semiconductive - 21-1-2026 at 21:58

It does have an odor.

But, there are confounding issues...

Amyl-alcohol is often described as fruity, maybe lime? But it stings my nose ... my nose isn't worth much ...

Under kerosene, the Amyl odor becomes faint.

I know the smell of sulfites in wine (allergies), and that's definitely not there. So, I don't think the sulfur dioxide is coming out in any significant amount.

I am not familiar with the smell of alkenes, like propene, so if methanol is being gassed I'm not sure what that should smell like.

I had to turn off the experiment (today) to solve some electronics incompatibility problems and back up software so as not to invite catastrophe. ( And, my sIster's coming to help me 'clean' tomorrow. uh-oh! The lab is in danger of being de-railed -- but her nose is better than mine. )

I noticed while it was off today that the solution rapidly darkened, and turned greenish (but not black).

It's as if being hot may have prevented a different side reaction from occurring.

I could re-run the experiment with no amyl-alcohol and just methanol next week, and then see if it smells like cookies. Is there a particular reason I should expect that smell?

I need to re-run the experiment anyway, to determine solubility of Na-M-bisulfite in 1,2,3-propanol. The words "very soluble" found on Wikipedia™ and everywhere aren't useful to me; I'm surprised chemists can do much with such sketchy data.

I'm not sure how to work out the solubility, efficiently.
I don't have a stirring mechanism in place for test tubes.

I'm thinking I'll just try to estimate how much might dissolve, and start at about 90% of that and add 5%, 2.5%, ... until it stops dissolving.

If I imagine that 1,2,3-propanol (glycerine) is water like, but with three oxy-hydrogen bonds instead of two -- then I suppose that 1 molecule of PPol is like 1+1/2 molecules of water in dissolving power. 2:3

I think 81.7 / 100 [g/cc] of sulfite is roughly how much dissolves in water near boiling:

So, I estimate:
Volume reduction of water from 20 [°C] to 90 [°C]:

100 [cc] · 0.9982055 ≈ 99.8 [g] of water = 5.04 [mol]

5.04 [mol-H₂O] · 2/3 [ PPol/H₂O ] · 92.09 [g-PPol/mol] ≈ 309 [g-ppol]

Therefore: ≈245 [cm³] of 1,2,3-PPol at 25[°C].
Does that seem reasonable?

If so: 81.7 [g] / 245 [cm³] ≈ 0.333 [g/cm³]

Plan of action: I measure out 1 [cc] of glycerol (and weigh it, to be more accurate), Then should expect around 333 [mg] of Na.M.BiS (total) to dissolve in PPol at 90 [°C].



[Edited on 22-1-2026 by semiconductive]

bnull - 22-1-2026 at 04:38

Quote:
I could re-run the experiment with no amyl-alcohol and just methanol next week, and then see if it smells like cookies. Is there a particular reason I should expect that smell?

About 15 years ago, when electrolysing an alkaline complex of Cu2+ in aqueous glycerol I noticed an odd smell, like cookies, coming out of the solution. I don't know what it was, apart from being an oxidation product of glycerol. Acrolein?

semiconductive - 22-1-2026 at 23:55

My sister smelled the test tube and concluded it smelled more like wine than cookies. So, I'll just go with that. Her nose is better than mine.

Did your glycerine turn yellowish (not that you could see it with Cu, I suppose) ?

I ask, because:
I have bought Schwan™ Glycerol USP, from the local shopping centers in the past. It came in a tiny/neat 40 [cc] or so sized brown bottle. It's meant to use as an emollient, or as a sweetner in confections (like cookies).

When I looked the data up online, I was informed that USP (grade?) could legally have up to 20% water in it, though the remainder had to be pure glycerol ... yada yada. The pages I read at the time suggested that it was very difficult to remove the water without damaging the Glycerol.

Several times I tried to raise the temperature of Glycerol to above the boiling point of water to 'dry' it. I didn't have precision temperature measurement at the time, so I can't tell you exactly how hot the bottom of the flask was; and I didn't know that kerosene or USP mineral oil were better floaters than limonene.

In air, glycerol would yellow extremely easily when any bubbles showed up. There would then be a strange odor that I would describe as almost burnt carmel candy.

A very similar smell would happen under limonene, at which point I just assumed that air wasn't needed to cause the yellowing and smell. I assumed that (maybe) just water is enough to decompose Glycerol -- or just heat.

Being disabled means I forget to press the enter key to start temperature regulation ... sometimes ... as in this morning with my sister talking to me...
And then an hour later realize that I forgot to plug the electrode back in after cleaning the desk up.

See where the Plot thickens:

pngsnap.png - 29kB

I'll smell it again in the morning.

bnull - 23-1-2026 at 12:13

I couldn't see it. The solution became a suspension of yellow to orange particles. The particles could be cuprous oxide or lead oxide, maybe even both, as I was using lead and graphite electrodes at the time. To make things more confusing, the lead electrode was an alloy of lead with a bit of tin that didn't darken or flake off excessively during electrolysis.

Edit: my glycerine doesn't discolor but fumes like mad and stinks of acrolein if I boil it.

[Edited on 23-1-2026 by bnull]

semiconductive - 24-1-2026 at 10:24

Some memories of cleaning up messes are probably best left in the past... :D

My 5 minute and 150 [°C] excursion, yesterday, did not noticeably change the odor of my test tube.
There's no yellow coloring either.

I'm mildly curious.
The temperature definitely went above the boiling point of water.

The temperature was not above the boiling point of glycerine ~290 [°C]
(I hope that data was discovered with a vacuum running!)

Acrolein:

If acrolein is created by hydrolysis, then less than 2% water (Duda™ Energy) might be too small to make acrolein; but (Schwan™) U.S.P. and <20% water might be enough water ?

On the other hand, if acrolein is primarily catalyzed by impurities; then I question whether Duda PPol vs. Schwan PPol has less of -- animal fatty acid, water, plant terpenols, ...

Then again, I used sodium metabisulfite, which is a mild reducing agent. eg: Sulfite might want to become sulfate before 1,2,3-PPol wants to become acrolein ?

To solve the puzzle:
I'll test straight Duda™ energy glycerol by cooking it in a test tube under kerosene. Let's see if it makes an odor by itself. I'll have to see if I can get Schwan™ at the store.

If I had a refractometer, I could estimate how much water that each PPol sample has. But, I just have a spectrometer; ( unfortunately numerical integrals done by Wolfram™ and other's, still dis-agree as of Jan-2026: https://physicsdiscussionforum.org/integration-of-planck-s-b... )

So, Let's see what I can do with density. I have bought a NEST lab pipette that transfers by using suction into a plastic tube. This isn't a highly accurate (or expensive) device, but I'm hoping it will last more than five uses which is all the disposable ones seem to handle before the plastic cracks.

I have a 1.000 to 5.000 [cm³] NEST device, and I bought two of them. If I can accurately measure 1 [cm³] of glycerol, then I might be able to figure out how much water there is from the mass.

Let's test with water first, which is more troublesome (vapor pressure is higher).

I'll set the NEST device to 1.000 [cm³]. Put 100 [mL] of reverse osmosis water in an Erlenmeyer in the microwave, and bring it to boiling.

Now I'll transfer water five times, into a weighed and tared test-tube.

!That didn't go well!

The erlenmeyer cooled down to 40 [°C] before the pipette was ready and assembled to use. I got between four and 5 samples before the water changed by two degrees [°C].

I had to learn not to push the button all the way down, but only to the point where it meets spring resistance. Otherwise it jams.

Note: The temperature recorded is the start temperature of the sampling:
Code:
T= 40.1 [°C] x̅= 1.186 σ= 0.044 [g] μ₀= 1.190 σ₀= 0.045 [g] 3.741% T= 34.9 [°C] x̅= 1.183 σ= 0.016 [g] μ₀= 1.185 σ₀= 0.016 [g] 1.346% T= 30.1 [°C] x̅= 1.176 σ= 0.012 [g] μ₀= 1.178 σ₀= 0.012 [g] 1.032% T= 25.1 [°C] x̅= 1.162 σ= 0.015 [g] μ₀= 1.164 σ₀= 0.015 [g] 1.326%


I got too much water each time and lots of variance. I'm not very consistent/steady at pipetting.

Let's try again!
I'll turn the dial down to 0.800 [cm³].

This time I'm aiming for 10 samples at each temperature.
I'll take the tip off, shake it out, and let fresh air in between every temperature run in order to make sure no water gets up into the suction chamber itself.
Code:
T= 70.0 [°C] x̅= 0.972 σ= 0.094 [g] μ₀= 0.988 σ₀= 0.096 [g] 9.703% T= 60.0 [°C] x̅= 0.985 σ= 0.030 [g] μ₀= 0.988 σ₀= 0.030 [g] 3.005% T= 50.0 [°C] x̅= 0.974 σ= 0.018 [g] μ₀= 0.976 σ₀= 0.018 [g] 1.893% T= 40.0 [°C] x̅= 0.963 σ= 0.019 [g] μ₀= 0.966 σ₀= 0.019 [g] 1.945% T= 40.0 [°C] x̅= 0.970 σ= 0.008 [g] μ₀= 0.971 σ₀= 0.008 [g] 0.781% T= 37.0 [°C] x̅= 0.961 σ= 0.013 [g] μ₀= 0.962 σ₀= 0.013 [g] 1.330% T= 35.0 [°C] x̅= 0.960 σ= 0.004 [g] μ₀= 0.960 σ₀= 0.004 [g] 0.368% T= 30.0 [°C] x̅= 0.965 σ= 0.006 [g] μ₀= 0.966 σ₀= 0.006 [g] 0.576% T= 25.0 [°C] x̅= 0.964 σ= 0.014 [g] μ₀= 0.964 σ₀= 0.014 [g] 1.492%


i took >90 samples, and the pipette is still working. Good enough.

I've computed the experimental statistics twice, first with standard math AKA: a Root of mean-squares algorithm; and second with math designed to estimate the mean of a population from a "Gaussian" sample, better. I take the square of a rooted mean.

I see I'm getting some decent repeatability ( sigma is small ) .... I just need more practice.

Uh... Wait... Water density goes down (less mass/volume) with higher temperature, doesn't it ?

I think this is what I'm supposed to be getting:

https://www.novabiomedical.com/education-training/knowledge-...

?! I get the opposite result !?


[Edited on 24-1-2026 by semiconductive]

semiconductive - 25-1-2026 at 21:36

If my Reverse osmosis water filter is expired, then I get some calcium through it in addition to CO₂ gas forming carbonic acid. But I expect briefly boiling would precipitate any calcium out -- so I think the water's volume vs. NIST specifications ought to be correct because I brought the water to a boil before hand.

If this pipette really is backward, then I suppose I could put pieces of nylon inside the pipette tip until I reduce the volume just enough that air expansion and temperature cancel out.

I can also put nylon washers under the depressor to stop it in the spring area at a repeatable spot. The only annoying weakness of this device is that the volume is set by spinning the depressor. It's not stiff enough to prevent me from accidentally changing it's volume setting when trying to rapidly use it.

Am I overlooking something?
Thoughts?

2026-01-24-16-04-53-153.png - 714kB

bnull - 26-1-2026 at 15:54

Quote: Originally posted by bnull  
Edit: my glycerine doesn't discolor but fumes like mad and stinks of acrolein if I boil it.

I guess I used a not so clean test tube. Today I heated some glycerine and, alas, it discolored to golden honey with smell of cookies. Ten years old double-distilled, pharmaceutical glycerine, 99.5% at least.

semiconductive - 28-1-2026 at 17:48

Fascinating. But, sometimes acrolein is clear...
I will probably get to cooking some 1,2,3-propanol in the next post.

I ran a whole bunch of calibration tests on my mass scale, today. Over 10 years, it's average drift is 6 [mg] out of 300 [g]. But, it annoyingly can thermally drift or get hysterisis by up to ±10 [mg] over an hour. ( It rarely does so, and only certain final masses, such as 'none' are followed by more than 100 [g] prior -- eg: that seems to be particular susceptible to it. )

Measuring my 50 gram standard is reasonably precise.

50.034, 50.034, 50.032, 50.030, 50.032, 50.033, 50.032, 50.031, 50.031, 50.029, 50.023 → μ=50.031 σ = 0.003 [g].

But, when I do it with 30 [g] + 20 [g] masses, I get: 49.996 σ=0.003 [g]
The same with smaller masses added together.
So, I think I probably need to steam the 50 [g] mass ... and/or (probably) replace it.

That's not a today job. But, it puts my pipette idea in perspective.

I'm confident that the scale is accurate to the same random bias error over it's whole range.

But: Still, the 1% variation that I get most of the time is pretty much within the scale's drift limit. It might not only be my pipette technique which is causing the deviations.

If I had my 3D printer running, I could print myself a liquid self-leveling mass scale that's accurate to 0.0001 [g]. But: Chicken egg problems ... I need to electroplate a steel rock tumbler drum to be 3D print accurately. eg: because I need iron free glass dust without moisture. And when I buy fumed silica from online, it's already been exposed to moisture and causes the resin to set even before I print it.

Unfortunately: Plating shops in portland, OR, refuse to do any work for the public since the government got into heavily inspecting them.

Not that tin is toxic, but what can you do...

I wonder if I can buy electroless nickel to make a protective coating over iron and if I need to sand blast the steel drum first...





[Edited on 29-1-2026 by semiconductive]

semiconductive - 30-1-2026 at 20:06

Cooked 1,2,3-Propanol. No color change up to 180 [°C] at the bottom of the test tube with 1.5 [cc] PP-ol under 5 [cc] kerosene (deodorized, super-pure). The thermometer at the top of PP-ol only registered 90 [°C] even after 20 hours.

Removing some kerosene, I got temperature up to 200 [°C] at the bottom of test tube and 101 [°C] at the kerosene-alcohol interface. However, 1 [cc] kerosene is not as good an oxygen barrier.

I decided to check what impurities, added, will do:

In order to get the PP-ol to discolor at all, I had to add some calcium chloride (anhydrous). That made the PP-ol turbid and a very slightly yellow. Then I added sodium-metabisulfite, and it floated on the surface (see picture) and caused gas bubbling. But it also made the slight yellow color begin to fade.

vlcsnap-2026-01-30-19h18m00s559.png - 307kB

I'm waiting to see if the bisulfite ever dissolves or not.
But, there's was never any odor.
Very curious.

Edit: The bisulfate began to dissolve, slowly, but as it did the 1,2,3-PPol turned black rapidly (less than 2 hours to totally opaque) and began to have strong odors. Duda™ energy, 1,2,3-PPol is advertised as 99.7+ pure. Mine has been opened for over a year, so I estimate it's still 98%+ still pure.

Calcium sulfite normally absorbs 4 water molecules to crystallize.

In order to test 1,2,3-PPol with a better oxygen barrier at higher temperatures, I'll need to use smaller samples of 1,2,3-PPol. Because it's a better thermal insulator than Kerosene is.

Redoing the experiment, 600 [mg] PP-ol, 2 [g] kerosene, unfortunately a couple of air bubbles got lodged in the PP-ol. I'll let it run anyway, and see at what temperature the bubbles dislodge.

The air bubbles never dislodged. They absorbed when the bottom of the test tube was at around 175 [°C]. This caused the PP-ol to golden-yellow slightly. I'm watching it to see if it will darken.

This photo is slightly misleading, because the camera is aimed along the refraction line. There isn't actually a dark band of material between the kerosene and the PP-ol. The color, to the eye, is golden yellow but not quite as dark as it shows in the photo.

vlcsnap-2026-01-31-16h58m31s915.png - 234kB

Note: The bottom of the test tube is holding steady at 195 [°C], while the kerosene just above the 1,2,3-PP-ol, is at about 125 [°C].


On this temperature graph, the air bubbles shrank between hours 3 and 5. This is where the yellowing occurred and the index of refraction of the 1,2,3-PPol changed remarkably. The white you see in the picture is a piece of teflon plastic behind the test tube so I could see the color changes of the liquid.

pngsnap.png - 34kB

My 30 watt soldering iron can't get any hotter, so this is pretty much the best I can do for a test tube experiment. I'm not sure if I can get a tiny heating mantle that would fit a test tube bottom.

There are temperature spikes in the graph, which is a bug in the INKBIRD bluetooth thermometer system. I'm working on that in another thread.

But, with a little bit more work I ought to be able to do time-lapse films of test tube reaction with text labeling on the video itself. :)


[Edited on 1-2-2026 by semiconductive]

semiconductive - 3-2-2026 at 16:17

Another thought occurs to me about obtaining sulfite ions.
How dangerous would it be.....

The solubility of SO₂ gas in water goes up as temperature goes down, and the gas doesn't all leave water when sulphite salts are acidified; on the other hand, there are other organics which actually complex SO₂.

According to Mellor in the volume on sulfur chemistry (see scimadness libeary), pp.~209,

Acetone complexes double it's own mass of SO₂ at 0 [°C].

If that's true, then I might be able to put sodium bisulfite or metabisulfite into the bottom of a test tube, cover it with 1[cc] of acetone mixed with ethyl citrate to reduce it's solubility in kerosene, and then cap the mix with kerosene. This would keep air out and (hopefully) prevent slow oxidation of sulfides into sulfates.

Then I can just allow dry HCl gas to bubble into the chilled solution at 0[°C].

The result should be table salt remains in the bottom of the test tube and acetone dissolves either H₂SO₃ or SO₂ under kerosene.

I should then be able to pipette the solution into a well stoppered Erlenmeyer, and put it in the freezer for short periods of time.

Also, Mellor pointed out something else that I find very interesting:

Mellor also notes, p.101, that K2CO₃ + S₈ can be dissolved in ethanol although neither is soluble in ethanol by itself. This suggests that potassium sulfite, or potassium polysulfides, may be formed by the reaction.

Since I suspect polysulfides might be the key to electroplating iron pyrite, so that's also worth a try!

semiconductive - 12-2-2026 at 23:00

Refraction data for water over temperature, how I am proceeding -- what I am learning:

I contacted someone in Turkey about an article on water refraction. In my last correspondence, they mentioned that they thought I was something like a 'data' assassin and recommended that I do a high precision refraction measurement myself and 'publish it somewhere' -- rather than try to figure stuff out from other people's data. ( So I won't mention her name here. )

I am attempting to follow her advice.
Building a refractometer:

I have capillary glass tubes, and am able to construct a refractometer based on the original Lorenz design:

https://riviste.fupress.net/index.php/subs/article/download/...

Jamin mirrors for optical surveying are common, and I have some.

I also have found that I can stretch glass fiber by melting it and pulling it, which thins the fiber and jacket. The shape becomes somewhat like a catenary/hyperbolic bend. This leaks more light depending on the refractive index of the fluid that surrounds it. So, I might be able to use a partially melted fiber as a second way to check the Lorenz measurements.

I don't have temperature control totally automated yet with calibrated glass thermocouple probes. ( I really need my 3D printer to work already!!! )

Tentatively,however, my first attempts match with the Turkish fiber optic experiment's plotted data. (which I can only get by picking it off a graph, as she doesn't publish the numbers!).

Which is to say, that my earlier comment to her on the her equation's dis-agreement with NIST has nothing to do with her actual data.
Her raw data isn't bad. BUT -- her published curve fit equation, and her data, do not describe the same curve.

I find the situation fascinating. She thought my pointing out the discrepancy was like an attack / discrediting smear campaign. Rather than as an opportunity to rise above her peers.

Things I have learned:

NIST's and IAPWS "data" isn't raw data.
Their data is actually a curve fit to the Lorenz-Lorentz equation.

The first article that I linked, ( Bashktov and Genina) from Russia, isn't actually using their own data (either) but is in fact curve fitting NIST's curve fit data. ( A curve fit of a curve fit !!! )

https://www.researchgate.net/publication/366494863_A_Simple_...


In all cases of disagreement that I've found, it is actually a curve fit equation that that is in disagreement with another curve fit equation.

It's not the 'raw' data itself. For all raw data I am finding is reasonably consistent (with random errors).

Therefore, I'm of the opinion that If the Turkish fiber optic experiment's raw data is correct, then the Lorenz Lorentz curve-fit equation makes slightly wrong predictions above 60 [°C]. I expect to get better data once I have 3D printing capabilities and a good calibrated thermometer.

But:
The 'upward' curve of the Lorenz-Lorentz equation in my earlier plots is actually evidence that a cube root of water volume will more accurately fit the data than multiplying by a raw volume.

Looking at Lorenz paper, it's fairly obvious that his derivation tries to match the leading terms of the Cauchy semi-empirical dispersion equation.

The same Cauchy semi-empirical equation is being used by both the various articles I linked as the 'simple' equation for their curve fits as well. I suspect something about the Cauchy semi-empirical equation is not quite correct.


[Edited on 13-2-2026 by semiconductive]

teodor - 13-2-2026 at 03:23

About your initial question of getting "free" (SO3)2- ion.
If we would skip a gass phase reactions that is possible only on condition of solution, that means the ion has mobility and can be separated from cation. There are 2 types of solutions: solid and liquid ones.
For liquid ones there are several methods. Some very sensitive acids could be produced by ion-exchange resins (e.g. you exchange Na+ with H+ or whatever). If you mean some space configuration, zeolites could be used to lock anions in specific cavities. But for electorchemistry I believe a common solution should be enough.
There are different solvents for SO2. All of them will give some type of mobile anion. Check SO2 solubility in different solvents.
As for water 1 kg at 0C it dissolves 3.55 mol of SO2 (Brasted, comprehensive inorganic chemistry, vol. 8). It is not bad. AcOH is far better solvent for SO2 by the way (also AcOH has well studied electochemistry).

If you plan to do some experiments you need not only know the method of getting free ions but some proof also that you have got it (because the experimental conditions always vary). For this purpose you have to start with pure SO2 and replace it with salts only when you get some qualitative results as the next step. The first step in chemistry is an experiments with pure substances and then you can replace it with more cheap ones or make the laboratory efforts cheaper but you need to get the reference data first. Without reference experimental data the cost of experiments is always high because of the initial step error impact. For this reason the zero step is literature search especially for experimental details - x g SO2, t=y C, got z % ions - repeat and check the experiment from literature, than start to vary conditions to make it closer to your schema. This is how you can proceed.

SO2 is not deadly toxic. It can be generated with Kipp's apparatus or just flask/dropping funnel (not so convenient as Kipp's because Kipp's produce only the amount required to saturate the solvent). Don't neglect a practical chemistry. You can easily store excess of SO2 in a freezer in a form of SO2*7H2O, those are nice harmless crystalls below 0C. Just add some water to get H2SO3 solution. As for more concentrated ionic solution SO3 or possible oleum would be an excelent solvent but it is much harder to work with.
In any case, I would just start with the solubility table of SO2.

P.S. As a practical obstacles for your final goals I would look into polymerization tendency of sulfur-oxygen acids, e.g. forming dithio acids which is how SO2 can behave in complex conditions.

P.P.S. You can't build anything by analysis. You can build something only by a synthesis. So, you use what you already have, know and tried to get more, know more and try more. There is no way to start with a mental construction and than implement it without having 99% parts of this schema aleady as a product of previous synthetic efforts. By this reason a practice of operating Kipp's apparatus is much more important to get some results with (SO3)2- ions.

As I noticed the chance to get a constructive response here is in reverse ratio to the length of the text in the message. Few people read something which is longer 2 or 3 paragrpahs at all. So, I intentionally try to respond to only one aspect of your question, getting free (SO3)2-, otherwise the discussion would be not maintainable. But I think I am over 2-3 paragraphs limit now ...

(Another aspect of starting a good discussion is to providing references to a literature. Your experimental details are for internal usage, and for discussion every result should be related to already known and published experiment. What was done before, what you did and what you observe as a difference. This way we can gain in a knowledge of a science. And this is not criticism, just my explanation of poor contribution from my side to potentially quite interesting thread).

[Edited on 13-2-2026 by teodor]

semiconductive - 13-2-2026 at 18:53

Quote:
There are different solvents for SO2. All of them will give some type of mobile anion. Check SO2 solubility in different solvents.
As for water 1 kg at 0C it dissolves 3.55 mol of SO2 (Brasted, comprehensive inorganic chemistry, vol. 8). It is not bad. AcOH is far better solvent for SO2 by the way (also AcOH has well studied electochemistry).


I am trying to.

I have glacial acetic acid (AcOH).
Also: According to Mellor (see two posts back) Acetone is an excellent solvent of SO₂.

My main problem is accurate equipment and methods for making measurements with in order to verify the results of what I've made.

My background is electronics, and if it's electrical I can make it work.
Chemistry -- I only took inorganic 200 level, undergrad.

I've been reading a lot of chemistry and physics publications;
But, I am finding mistakes in articles that are quite annoying and even which undermine the credibility of experiments that I can do.

For example, Lorenz Lorentz theory was derived using a semi-empirical Claussius Messoret formula. Both researchers used the same crude approximation, which is isn't for molecules but for a homogenous cavity.

To compensate for hetrogenous molecules in a liquid, giving different refraction indicies, requires using a discrete electronic model ( and at least a reference to the Vanderwall's equation of state for liquids and gasses near boiling! Which I can do. ) But, neither of these was done in the case of Water even up to the boiling point by NIST. So, NIST's article was actually misleading to me.

I only discovered that the LL equation was approximate after reading and doing some computations based on the suggestions of D.E. Aspres, Bell Laboratories in American Journal of Physics, Vol 50, No. 8, Aug 1982.

https://www.researchgate.net/publication/235409855_Local_Fie...
Quote:
SO2 is not deadly toxic. It can be generated with Kipp's apparatus or just flask/dropping funnel (not so convenient as Kipp's because Kipp's produce only the amount required to saturate the solvent)


I have an addition funnel, formic acid 95%, and sodium meta bisulfite. I've been told that will generate SO₂ gas by several chem sites on the internet. I was thinking of setting up a dripper.

But: I was dubious because formic acid isn't actually a strong acid; even citric acid is stronger than formic. I tried formic acid under kerosene and added Na-MBS: Not surprisingly, I didn't get any gas bubbles passing through the kerosene.

Given what you say about AC-OH, I think the odds that formic can dissolve SO₂ are probably high. There is only one saturated carbon atom difference, after all.

There is so much (unqualified) information on the web without conditions recorded, that it's difficult for me to design an experiment that has a reasonable chance of succeeding.

Note:
The experiments I just did, in the last 10 posts, essentially demonstrates that glycerol under kerosene (anoxic conditions) does not yellow up to 150 [°C]. The yellowing I'm getting, apparently is an aldehyde. For aldehydes form adducts with sodium meta bisulfite, and the yellowing faded when I added Na-MBS. Alcohols don't form adducts with Na-MBS.

So, I know enough about glycerol (now) that I am reasonably confident that if I added SO₂ gas to it, that I could at least tell if the gas was dissolving or bubbling out.

Note:
I also have iron pyrite, I can get muriatic acid, and that will produce SO₂ gas in bulk. I'm not sure how to keep HCl gasses and moisture out of the product -- but I do have CaCl₂ and a drying tube.

I don't have a kipp system yet -- but I also have test tubes with gas outlets near the bottom that can be stoppered. I could connect two of them together with silicone hose to generate SO₂ in one test tube, and dissolve it in the neighboring test tube in acetic, or 1,2,3-propanol, or methonol, etc.

bnull - 14-2-2026 at 16:26

Quote:
I also have iron pyrite, I can get muriatic acid, and that will produce SO₂ gas in bulk.

It won't work. Iron pyrite is iron sulfide, not sulfite.

semiconductive - 15-2-2026 at 14:56

Quote: Originally posted by bnull  
Quote:
I also have iron pyrite, I can get muriatic acid, and that will produce SO₂ gas in bulk.

It won't work. Iron pyrite is iron sulfide, not sulfite.


Doh! H₂S. I knew that.
Yes, muriatic acid + na-MBS, then.

teodor - 16-2-2026 at 15:53

I’ve attached a chapter from Seidel / Linke book, it is about SO2 solubility in different substances.

My idea is that because any solubility which is different from what is predicted Raoult’s law means some chemical reaction between solvent and solute (a solvate formation) and in many solvents that solvate is charged (but not always as SO3- ion, and I assume that is not a strict requirement) the solvents which dissolve SO2 better are more capable of the solvate formation.

We can also probably predict that solvate is charged in those solved which are self-ionizing. This is not limited to prototropic solvents (water, ammonia, anhydrous acids). Liquid SO2 itself is oxidotropic solvent and it dissociates to ions:

SO2 <-> SO++ + SO3—

Some salts like KI or KSCN are very soluble in SO2 and can elevate the boiling point, I assume up to a room temperature. This is also could be true for compounds like toluene, mixture of nitrobenzene and SO2 and others which have high solubility numbers or low vapour pressure of the system.

Also at some point with liquids a solution of SO2 in a solvent can become solution of a solvent in SO2.

Which results you would like to measure and what is the equipment you try to use?

Formic acid is stronger than citric acid. And also stronger than acetic acid. But I assume it is not so convenient to work with.

SO2 doesn’t react with O2, so if you need an oxygen scavenger (you mentioned anoxic conditions) it can not serve the purpose.






Attachment: SO2.A.pdf (1.8MB)
This file has been downloaded 12 times

semiconductive - 18-2-2026 at 16:22

Hmmm...
When I look up pka₁ of citric acid, I get 3.13, pka₂ is 4.76. etc.

? Don't Larger numbers means a more tightly bound H atom and a weaker acid ?

When I look up formic acid, I get pka₁ = 3.87.

3.13 < 3.87, so I am under the impression that for first ionizations (infinite dilution) -- citric is stronger than formic.

I have 95% formic, and 99.9+% glacial acetic. Either is fine to use. Although my formic is old enough that it's slightly discolored due to light induced chemical reaction or oxidation.

Unfortunately, a mouse broke my glass HCl jar (distilled, 40+% using a vigereaux and gas trap) and it predictably exploded into a cloud of gas -- ( but the mouse survived!!!) . So, I'll need to get muriatic at the hardware store and re-distill it, or use just use sulfuric acid.

Because my vacuum pipettes suck for accuracy ... I also bought glass borosilicate 3.3 TD 1 ml pipettes designed for 20 [°C]. It has 0.01 [cc] gradulations. I was going to buy some temperature regulators for my milligram scale, and see if I can get the precision of my scale to increase. These are the only tools beside reagant grade chemicals that I presently have to make measurements with that could possibly make accurate denisty measurements to figure out anything about dissolved gasses or ions.

I have an optical spectrometer, but I don't have a way to calibrate it yet.
I also have multiple thermometers, but they are only partially functional. ( I'm paying someone to fix it, but they are slow. )

----

Roults law shows that the boiling point (AKA out-gassing point for SO₂?) is affected by the molar fraction of each substance in solution. This is similar to Henry's law.

These require solubility calculations, etc.

The Russian article I linked to earlier, has a few example solubility calculations that ought to be parallel to the SO₂ gas calculations I'll need to do.

These calculations should be related to Raoult's law and Henry's law.
So, let me give an example of what I know and where I get stuck:

I was taught in undergrad chemistry that mole fraction calculations are done by writing reactants over products, and dropping any pure water, electron, or solid terms. The purpose of my chemistry classes was just to teach me how to read literature and have a basic understanding of how the calculations are done so I can follow other people's instructions.

Here's an explanitory article for just water ionization, and solvation of gasses in water:

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

Equilibrium equations (2) and (3), talk about diatomic gas molecules dissolved in water.
They aren't clear exactly what state the gas is in, and because there are electron transfers I assume you will treat these equilibriums as a chemical reaction and not just a dissolution process:

H₃O + 2e⁻ ⇌ H₂ + 2·H₂O #(2)
1/2·O₂ + H₂O + 2e⁻ ⇌ 2·OH⁻ #(3)

If I assume the gas is the reactant, and the ions are the product, then I will get a specific constant that tells how hard it is for a diatomic gas to turn into ions:

KH₂(aq) = [ H₃⁺ ]² · [2e⁻] / ( [ H₂ ] · [H₂O] )
KH₂(aq) = [ H₃⁺ ]² / ( [ H₂ ] )

Therefore I get:
[H₃O⁺] = √( KH₂ · [H₂] )

Which agrees with the Russian paper's stated equation on p. 244.

However, If I do the same with the diatomic oxygen ...:
1/2·O₂ + H₂O + 2e⁻ ⇌ 2·OH⁻ #(3)
O₂ + 2·H₂O + 4e⁻ ⇌ 4·OH⁻ #(3)

KO₂ = ( [ OH⁻ ]⁴ ) / ( [H₂O]² + [ O₂ ] + [ e⁻ ]² )
KO₂ = [ OH⁻ ]⁴ / [ O₂ ]

I get a fourth root:
[ OH⁻ ] = ∜( KO₂ · [ O₂ ] )

Which is not the same as the Russian article's [ OH⁻ ] = 2·√( KO₂ · [ O₂ ] )
And I don't follow how they got the equation that they got.

Later in the paper, they talk about non-ionized dissolved gas in water.
But, they went the route of a diatomic molecule dis-associating in those calculations, and I've never seen that kind of calculation in U.S. literature. ( My chem class is worthless in understanding what I'm supposed to do. )

I'll need to do the same for SO₂ in solution, to determine equilibriums, after reading your paper carefully.

But, there's SO₂ gas (which I assume is a gas bubble in solution, and which has a volume of a gas per molecule); then there's SO₂ (aq) which is hydrated with at least one water molecule and doesn't obey the gas law for volume any more, then there are the ionized states of SO₂.

And, I'm not sure how to treat each one in terms of molar equation writing and constants.
More to come (below) after I think about the paper you gave me....

semiconductive - 18-2-2026 at 19:41

Thank you for the paper, Teodor.

From reading it, I notice that there is a way using an iodine solution to titrate SO₂, in order to figure out quantitatively how much is dissolved. Do you happen to know what kind of solution is used and how? That seems like something worth ordering and keeping on hand in my lab.
Quote:
SO2 doesn’t react with O2, so if you need an oxygen scavenger (you mentioned anoxic conditions) it can not serve the purpose.


This brings up one of many issues that are partially confounded and not recorded in the data ( including the book chapters you just gave me to read. )

Water, upon standing, will absorb diatomic oxygen and nitrogen from the air.

Yet, the experimenters do not say what condition the water was prepared from nor how long their experiments took to complete, nor what precautions were used to exclude atmospheric air from the experiments.

Often, I have found (to my dismay) that lack of recording conditions may mean an experimenter never even thought of the issue and that their data is contaminated with unknown variables.

From my research, I think it's generally sufficient to bring distilled water to just 99.9 [°C], in order to expel air molecules from it. But, something has to be done in order to prevent the cooling water from re-absorbing air and CO₂ immediately. Freshly distilled water is naturally an-oxic, although the Russian paper that I read suggests that it may generate it's own oxygen because of chemical equilibrium issues. I don't know on what time-scale to expect oxygen to spontaneously appear in water.

For now:

I simply place kerosene (ultra pure) on top of my liquid mixtures. O₂ is more soluble in cooling kerosene than in water, so kerosene acts as a sponge and slows the penetration of O₂ into a freshly boiled solution of water. It's also possible to get powdered aluminum or tin and place it in the kerosene when in a glass jar, and then expose it to UV light in order to initiate reactions with oxygen to remove the dissolved moelcules.

I also have CO₂ on hand and can use that to displace air above kerosene. (Although I've been lazy, so far... ) CO₂ is polar, and pretty much won't dissolve in kerosene.

I'm open to suggestions, if you have any, for other ways to limit the access of oxygen to my experiments. :cool:

But part of the issue, is I don't have a way to measure how much oxygen has penetrated into my solutions until something obvious (like rust) shows up.

semiconductive - 18-2-2026 at 19:48

Note: Sulfite ions,are mentioned in the literature as absorbing and reacting with oxygen (in what form, I don't know) -- and transforming slowly into sulfuric acid. This is an un-desirable side reaction which I want to avoid.




[Edited on 19-2-2026 by semiconductive]

teodor - 19-2-2026 at 15:03

Hm, I think the solution is some salt of a transitional method with lower valency, like Cu(I), V(II) (extremely powerful scavenger) etc. But e.g. Zn will not work - it will form dithionite with SO2, which is also very fast O2 scavenger, but it reacts with O2 to form SO4-- ion which possible you try to avoid.

Oxygen can easily penetrate any liquid which can dissolve it. The lowest solubility is in hexane, paraffine oil and highest alcohol (C10+). Kerosene is a mixture of hydrocarbons and I think oxygen has good solubility in it. You need only pure saturated hydrocarbons like medical paraffine oil. Or pure hexane/heptane etc.

And indeed, over time Na2SO3 solution is converting to Na2SO4 even in closed bottles.

It is better operate in a closed system allowing scavenger to destroy all the existing oxygen before start of the operation.

You also mentioned that solubility in SO2 is by forming "molecular cavity" if I understood your previous messages correctly. It is not always so and highly dependend on the solute.

[Edited on 19-2-2026 by teodor]

semiconductive - 19-2-2026 at 17:55

This post is a derivation of an improved model for refraction index of a substance vs. density.

Relative dielectric constant of a substance is a value equal to the index of refraction (of that same substance) squared. It varies with frequency, but is constant at any specific frequency.

ε_r = n²

Which means that a Lorenz-Lorentz model OUGHT (but isn't) to be identical to a model where every molecule is an electronic capacitor having a constant dielectric constant inside and the space between molecules is a capacitor with the dielectric constant of empty space.

Unfortunately Lorenz and Lorentz, both made a simplifying assumption that is not rigorously correct in three dimensional volumes of liquids or gas. They used a polarization model with a single charge cavity and no corrections for series vs parallel dipole moments.

eg: This neglects the different series and parallel effects of molecules spread out in space. Capacitors do not add in only one way.

For an analytical solution that predicts the quality of the error made by Lorenz and Lorentz, I'm going to make a simple linear assumption. I am going to assume that molecules can be modeled linearly as a cubic shaped capacitor with two imaginary conducting plates oriented perpendicular to one arbitrary dimension of the space.

For water, then, the mass and volume of each water molecule 'capacitor' can be found from the maximum water density 3.98 [°C].

Note: At 980 [nm], 3.98 [°C], 1 [ atm ] the refractive index of water is computed as 1.3266 from IAPWS formula. This means the dielectric constant is ε_r = 1.3266² at 980 [nm].

If the total volume of water material is (1 [cm])³, but the total volume we measure the capacitance of is is (B [cm])³ , then the difference in volumes is empty space between and around the water molecules.

Note: It does not matter how many smaller molecular cubes the 1 [cm] cube of water is broken up into, because in electronic circuit theory, capacitors are linear circuit elements. The only things that *really* matter are the relative number of water molecules found longitudinally ( in series ) when passing through any given cube vs. transversely ( in parallel ). For this determines the number of series to parallel elements in a capacitor network.

But, this means I can compute the exact same (bulk) capacitance answer by imagining all water swept into a corner of the volume B³ and the rest left as empty space.

By definition, the density of the entire cube is still 1/B³ even though we've compressed the 1 [cc] of effectively 3.98 [°C] water into a single corner of the cube.

The capacitance of the entire cube can be computed from a network of cuboid shaped capacitors that contain one and only one kind of 'thing' each.

I break the entire cubic volume (B³) up into one cuboid capacitor formed of maximum density water, and three capacitors of different sizes that represent the remaining (empty) cuboid spaces around the water.

The nominal capacitance of a cuboid is defined as:
C = dielectric constant · area_of_face / length_between_faces

I can then describe the entire volume as four cuboid capacitors.
C₀ = ( ε₀ · ε_r ) · 1² [cm²] / 1 [cm] # Capacitor made only of water
C₁ = ε₀ · 1² [cm²] / (B-1) [cm] # Empty space in series with water
C₂ = ε₀ · (1·(B-1)) [cm²] / (B-1 [cm]) # Small space in parallel with water
C₂ = ε₀ · (B-1)·(B) [cm²] / (B [cm] ) # Larger space in parallel with water

I'm neglecting electric field fringing effects, and molecule shape effects, because the error is a fixed percentage based on geometry and this error usually cancels when converting a capacitor value back into a dielectric value. The capacitance values are crude, but the dielectric values ought to be quite accurate.

The total capacitance of the composite cube is:
C = 1/ ( 1/C₀ + 1/C₁ ) + C₂ + C₃
C = ε₀·( 1/( 1/ε_r + B - 1) + 1-1/B + B-1 )

Therefore, for the whole cube, the average dielectric constant is:
n² = ε_a = 1/( B² - B + B/ε_r) + 1 - 1/B²

eg: This is (theoretically) the "average" or "equivalent" dielectric value that light would encounter at a given density when moving through equally dispersed water molecules in a volume.

---------------------------------------------------------------------

Here is a graph of IAPWS prediction for 980 [nm] light's index of refraction in freshly distilled water from freezing to boiling at 1 [atm]. I've also included a plot of what my lumped electronic circuit model predicts for the same color of light.

pngsnap.png - 14kB
Code:
#!/bin/env gnuplot # Compute an 'average' index of refraction for a 1 [cm³] volume of ε_k media # distributed in a vaucuum volume of B³, where refractivity is ε₀. # Written by Andrew Robinson of Scappoose, 2026 # Copyright 2026; Released under the GNU pubic license GPL3.0. # https://www.gnu.org/licenses/gpl-3.0.html # # 980 [nm], IAPWS calibration n_k = 1.3266 E_k = n_k**2 # Block model of how permittivity changes with molecule spacing E_B( B ) = 1/( B**2 - B + B/E_k ) + 1 - 1/B**2 n_d( d ) = E_B( (d/1000.)**-0.3333333333 )**0.5 # IAPWS -- characteristic equation # l is wavelength in nanometers # p is density in kg/m³ # T is temperature in kelvin R997(l,p,T)=(_a0+_a1*(p/p0)+_a2*(T/T0)+_a3*(l/l0)**2*(T/T0)+_a4/((l/l0)**2)+_a5/((l/l0)**2-Luv**2)+_a6/((l/l0)**2-Lir**2) + _a7*(p/p0)**2 )*(p/p0) _R997(x)=sqrt( (1+2*x)/(1-x) ) # Caution, this inputs temperature in Celsius and converts it to Kelvin. R997N(l,p,t)=_R997( R997(l,p,t+273.15) ) T=293.15; T0=273.15 l0=589.0 Lir=5.432937; Luv=0.229202 p=997.0 ; p0=1000.0 _a0=0.244257733 _a1=0.00974634476 _a2=-.00373234996 _a3=0.000268678472 _a4=0.00158920570 _a5=0.00245934259 _a6=0.900704920 _a7=-.0166626219 # IAPWS -- density equation for water 0-100C ( <0.02% error ) # https://chem-casts.com/knowledge/density-of-water p(T) = ( 999.83952 \ +16.945176*T \ -7.9870401e-3*T**2 \ -4.6170461e-5*T**3 \ -2.805425e-10*T**5 )/(1 + 1.6879851e-2*T) # set title "Index of refraction at 980 [nm] for freshly distilled water vs. temperature" set dummy t set grid set xlabel "Temperature [ °C ]" set xrange [ 0 : 100.0 ] set xtics 5 set yrange [1.307:1.330] set ytics auto set mytics 5 set ylabel "Refractive index, n" plot R997N( 980.0, p(t), t ) lw 3 lc 'grey', n_d( p(t) ) ti 'block dielectric'


What is important in this graph is not how accurate the IAPWS density vs temperature equation is, but the relative shape difference between the IAPWS curve and my own prediction of index of refraction based on density.

The IAPWS graph represents assumptions based on a Cauchy curve fit of the index of refraction data. My curve has a very similar shape but is slightly higher in value in general but crosses over near boiling.

Now:
Find a graph of experimental data for water at 980 [nm] using fiber optics, vs. a Cauchy curve fit in an article referenced in one of my earlier posts. (Esra K. and and S¸ Yaltkaya are authors, Indian Journal of Physics. ) Esra's actual fiber optic data at 980 [nm] tends to be above her Cauchy curve fit.... (Equation #5 in the article) and crosses over near boiling.

eg: The actual data compared to a Cauchy curve fit agrees (qualitatively) with the electronic correction model I just computed compared against a standard IAWPS curve fit.


[Edited on 20-2-2026 by semiconductive]

semiconductive - 19-2-2026 at 20:55

Quote:
Oxygen can easily penetrate any liquid which can dissolve it. The lowest solubility is in hexane, paraffine oil and highest alcohol (C10+). Kerosene is a mixture of hydrocarbons and I think oxygen has good solubility in it. You need only pure saturated hydrocarbons like medical paraffine oil. Or pure hexane/heptane etc.


Ultra pure lamplighter's kerosene fuel, on the label, says made from paraffin.
Quote:
Cu(I),


I know how to reduce copper chloride using ascorbic acid, but I don't think it oxidizes easily afterward. I'll have to look into that.
Quote:
You also mentioned that solubility in SO2 is by forming "molecular cavity" if I understood your previous messages correctly. It is not always so and highly dependend on the solute.


Water is not a very homogenous substance.

Avagadro's number was estimated by Lorenz (a dutch physicist) based on refraction inferences; but his approximation is a homogenous one plane calculation and has the correct exponent ·10¹⁹ but isn't even one figure accurate in the decimal places.

Yet, his equation is what IAPWS is using for calculating water density and states.

The equation I derived in the immediately preceeding post is a modern correction to the L-L refraction formula. But, the main difference is that it assumes a hetrogenous matrix of molecules such that light can move alternately through vacuum and through water either in parallel or series.

It is not a 'homogenous' single cavity approximation.

Anyplace around a water molecule that isn't 1 gram/cc of space density, must be modeled as a cavity of vacuum that the molecule is inside of.

In more modern theory, vander-walls computes a 'volume' or radius of distance between molecules that they 'will not approach' each other.

These effective 'volume' ideas are critical to computing chemical reactions in liquids.

But this vanderwall's radius is a statistical average of some sort.
Because, eg: water is not all at exactly the same temperature or pressure (especially when in a container, and gravity pulling on it!)

If the average temperature of water is 99 [°C], then we might estimate that a bell curve describes how much of the water is hotter or cooler. ( or another curve ).

But the distributed nature of water means there will be a very small number of micro boiled places which are effectively at 101 [°C] in a bath of water which measures 99 [°C] with a thermometer no matter where you place it.

The 'state' of water is not described, exactly, by a single temperature.

Some water is in liquid form, and some of itcan be in gas form (This may be true even at freezing, although the percentage of water at boiling with the average water temperature at freezing might be less than 1 part in a trillion...).

The magic of the mass-action law is hidden inside statistical averages.

The beauty of the vanderwalls approach is that gas and liquids are not separate states, but that a single gas law can describe both liquid and gasseous states.

However, (obviously) the total density of water (or any liquid) will depend on how much of it is in gas form vs. hydrated form, vs. double-hydrated form, etc.

Is is the hetrogenous nature of water which pH is describing, and that's not the only hetrogenous effect present.

This is what was so disturbing to me when reading the Russian article. I naively thought that either hydrogen gas was a gas in water, or it chemically combined to make hydronium ions.


But the Russian view doesn't just accept that gas H₂ dissolves in water. He also assumed that there is disassociated and un-ionized hydrogen dissolved in water. Therefore, we have four distinct species [ H₃O ], [ H₃O⁺ ], [ OH⁻ ], and [OH] in even pure water; and each will have a different vander-walls radus as a liquid.

And, now, I realize that I probably also have to consider H₂ gas as a distinct thing taking up a different volume from any of the previously mentioned ions.

So, I'm looking at SO₂ with trepidation, and thinking -- how many ways can this thing become distinct?

The relationships must all be governed by the mass action laws; it's just a matter of accounting for all statistically significant possibilities and figuring out equilibrium constants.

bnull - 21-2-2026 at 16:08

Quote:
I know how to reduce copper chloride using ascorbic acid, but I don't think it oxidizes easily afterward.

Cuprous chloride is easy to oxidize and is photosensitive, especially when wet.

semiconductive - 21-2-2026 at 16:19

I know that cooking glycerol turns it a golden yellow when exposed to oxygen at temperatures above 155 [°C].

I don't have a way to calibrate my spectrometer for photon counts; but I certainly could compare the relative brightness of two color pixels after adding known volumes of oxygen to 1 gram of 1,2,3,PP-ol and cooking it. This would give me a way to correlate amount of oxygen exposure in the form of air bubbles absorbed by glycerol to color change.

That, at least, would allow me to objective test how much ultra-pure kerosene leaks oxygen vs. paraffin. (I have solid USP paraffin as well)

eg: I could watch how long it takes for glycerol to yellow underneath paraffin wax vs. lamplighters kerosene.

On rather curious point, ultra pure kerosene can be frustrating in that when the liquid below it has been boiled that it will suddenly begin mixing with the upper layers. It's as if the polar vs. non-polar idea stops working and the two liquids mix once boiled. Only density makes a difference as to which liquid floats and which one sinks.

But, I'm reasonably confident that it excludes oxygen better than most other organic liquids that I've tried. Hexane isn't one I've tried because it's listed as a mild neurotoxin.

semiconductive - 21-2-2026 at 16:46

Quote:
Cuprous chloride is easy to oxidize and is photosensitive, especially when wet.


So, I mix cupric chloride with ascorbic acid, and I will get cuprous chloride . If I'm recalling correctly, the Cu-(I)-Cl will precipitate as a slightly off white powder.

Question:
I can certainly run the precipitate through a 15ml buchner filter and rinse with distilled water. Would this be enough to get rid of the ascorbic acid?




teodor - 23-2-2026 at 02:05

Quote: Originally posted by semiconductive  
Hmmm...
When I look up pka₁ of citric acid, I get 3.13, pka₂ is 4.76. etc.

? Don't Larger numbers means a more tightly bound H atom and a weaker acid ?

When I look up formic acid, I get pka₁ = 3.87.

3.13 < 3.87, so I am under the impression that for first ionizations (infinite dilution) -- citric is stronger than formic.

I have 95% formic, and 99.9+% glacial acetic. Either is fine to use. Although my formic is old enough that it's slightly discolored due to light induced chemical reaction or oxidation.



Do you have the conditions of infinite dilution in water in your experiment? pKa is solvent-dependent. When we talk about absolute acids, I don't want to have any formic acid on my skin.
Formic acid can react with alcohol to form ester, citric acid is not. So, I assume their strength in non-aqueous conditions are reverted.
Of course, absolute acetic acid is a better solvent than formic acid from practical point of view because it is stable in pure state and formic acid is not.

But I agree, I have no right to talk about acid strength without precisely understanding the conditions. Just from a practical point of view, don't treat a formic acid as a weak acid.

Also I am still reading your others comments when I have time, so may be I will respond to some other your thoughts. But there is general remark: there is no one theory of solubility. So, there is no one generic model you can use in all situations. That's why those solubility data get my measurement is important. There are many ways how water can interract with SO2 and even more ways how SO2 can interract with different substances and not all those ways are well-studied yet. Some of them are studied. But there is no one general formula or theory which can be used in all those cases.

[Edited on 23-2-2026 by teodor]
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