Sciencemadness Discussion Board - Production of Cu (I) Cl with household chemicals. Problems.

But I am curious about what drives the metabisulfite reaction at all, or how sodium gets stripped from chlorine given the pka of HCl is -7, but that of SO4 is -3 or +2.

The sodium isn't "stripped" from anything- the sodium is an ion, and is only hanging around the chloride or sulphate because they've got opposite charges. Once it's dissolved in aqueous solution, the sodium ion is solvated by water, and couldn't care less about the anions.

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In your suggestion, the sodium is stripped by the sulfate to make sodium sulfate and copper II chloride. So, there is some sodium to neutralize the bisulfite, but only if the Sulfate ion gives up the sodium. Could you explain why the experiment would still work with the PH of the solution going very acidic? I mean, will copper I chloride remain a precipitate with free sulfuric acid around? or will some ppt re-dissolve into solution?

No- the sulphite ion can reduce the copper(II) to copper(I) only when there are chloride ions around to stabilize copper(I) as the insoluble chloride. The sodium ions aren't involved in the reaction at all, and are only there because you can't add chloride without some cation along with it. The pH for this reaction scarcely matters because chloride is a very weak base, and H+ ions will not interfere with the precipitation of the CuCl.

[Edited on 29-11-2017 by DraconicAcid]

semiconductive - 29-11-2017 at 00:44

Take your green hydroxide/chloride and dissolve it in hydrochloric acid (muriatic acid), add some electrical-grade copper wire, and seal it. The solution will be yellow if all of the copper is in the +2 state, and brown-black if it's a mixture of +1 and +2. Shake it with the wire until the solution goes colourless- this may take overnight.

Dilute the solution by rapid addition of boiled water, and white CuCl should precipitate.

OK. That's similar to the idea I was told to do on the other thread.
The only difference is that they told me to boil them all together, and you're telling me to first let them sit until it's colorless and then add boiling water.

In the experiment I tried, I put copper wire, salt, and copper sulfate in a jar. There was a stoichometric mix of salt and copper sulfate. The color of the solution went green immediately. The cap temperature was set at 40 degrees, for several hours. During that time, the copper wire was coated with a precipitate that could have been white or light green and no other change to the greenish solution happened. (It's hard to tell copper wire color inside green solution). At the end of several hours, I began raising the temperature to see if the copper wire would clear up ... and it did as I got above 70C. By 80C, the wire was clean. As I approached boiling, the green precipitate began forming (described in other thread). https://www.sciencemadness.org/whisper/viewthread.php?tid=62...

So, my solution was always greenish. Neither yellow nor brown black were ever seen.
Is this due to a contaminant in the morton salt, KI, or the anti-caking agent?

1st: Given your process, will it work with stoichiometric 2 NaCl + CuSO4 --> Na2SO4 + CuCl2 , or does there have to be excess HCl for it to work, and why?

2nd: What is the maximum temperature that the solution ought to be kept at overnight to dissolve the wire?

[Edited on 29-11-2017 by semiconductive]

Boffis - 29-11-2017 at 01:14

If you have time the simplest method is to place bare copper wire or bits of copper in a jar and fill it up with dilute ammonium chloride solution and seal tightly. Cuprous chloride crystals grow slowly as magnificent complex modified cubic crystals to 8 mm or more, though it take years to get large ones. I discovered this process by accident but it yields very nice crystals which oxidize more slow on exposure than the powdered form. When I opened my jar after many years I found that there was no smell of ammonia which I find rather curious.

DraconicAcid - 29-11-2017 at 08:15

So, my solution was always greenish. Neither yellow nor brown black were ever seen.

That's because you had a lower concentration of chloride ion.

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1st: Given your process, will it work with stoichiometric 2 NaCl + CuSO4 --> Na2SO4 + CuCl2 , or does there have to be excess HCl for it to work, and why?

Higher concentrations of chloride help keep copper(I) in solution so that it doesn't coat the wire. Wire won't react if it's coated.

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2nd: What is the maximum temperature that the solution ought to be kept at overnight to dissolve the wire?

I don't think it matters, as long as HCl isn't evaporating.

[Edited on 29-11-2017 by DraconicAcid]

JJay - 29-11-2017 at 08:55

Quote: Originally posted by Chemetix

Copper (II) sulphate + NaCl => Cu(II)Cl2

Cu(II) Cl2 + sodium metabisulphite => Cu(I)Cl ppt

Decant and rinse and decant again.

Total bucket chemistry in a one pot synth.

Beautiful.

semiconductive - 29-11-2017 at 12:38

Quote: Originally posted by Boffis

If you have time the simplest method is to place bare copper wire or bits of copper in a jar and fill it up with dilute ammonium chloride solution and seal tightly. Cuprous chloride crystals grow slowly ... When I opened my jar after many years I found that there was no smell of ammonia which I find rather curious.

I'm buying 2 dram vials with phenolic screw on caps so they can be heated up to 150 to 200C under pressure. It's no biggie to leave one of those with the solution in it on an unheated rack and see what happens. I'm more concerned with safely producing ammonium chloride without making ammonia and chlorine gas... probably something to do outside!

I don't think it's all that strange you didn't smell ammonia. I made a copper oxide powder (black), and placed it in aqeous ammonia (15%) with oil over it to prevent air from getting in. The ammonia did not dissolve the copper oxide, as far as I could tell, even after a week of sitting. But, after removing the oil the same bottle exposed to air began forming an indigo colored solution within minutes especially if I blew across the water. What is going on is that CO2 from the air attacks the copper oxide and the aqeuous ammonia; producing ammonium carbonate and possibly copper carbonate. The ammonia odor becomes very faint over time, as complexed ammonia with copper is more stable than ammonia carbonate by itself.

I would think that the ammonia chloride in your solution would also make a complex with copper chloride, as it forms. So, there ought to be both copper chloride (hydroxide?) held in solution by ammonia, and your precipitated crystals.

Did you weigh the remaining solution to get any quantitative idea of how much dissolved material was still in solution after the crystals precipitated?

I'm pretty disgusted with my pyrex volumetric flasks. Before I do your experiment, I am going to flame polish the bottom and then see if I can deform the concave portion of the bottom to correct the -0.200ml error in volume that I measured. The flask says +-0.05ml, but I'm getting a lot bigger error than marked. If you have any suggestions for how to accurately deform glass ... I'll be thinking about ways to do that tonight.

[Edited on 29-11-2017 by semiconductive]

semiconductive - 29-11-2017 at 16:50

But I am curious about what drives the metabisulfite reaction at all, or how sodium gets stripped from chlorine given the pka of HCl is -7, but that of SO4 is -3 or +2.

I'm after an understanding of the principles involved, not just making chemicals. I'm wanting to know why my experiment didn't work, as well as why yours will.

Charges don't have brains, or "care"; I get that. However, I've been taught that semiconductor physics is analogous to aqueous chemistry. I originally took chemistry as an EE. The idea my chem teacher gave is a little more complicated than what you seem to be saying.

In semiconductor physics, free wandering "ions" in the crystal are attracted to the opposite charge just like ions in water. There are "bound" states, where an ion (fermion) hangs around the oppositely charged ion but never wandering more than a few angstroms away; eg: in reality the "bound" ion is actually "dissolved" in the semiconductor crystal in a molecular orbital. It typically only wanders around one to two atoms away from the ion that it is actually attracted to. When the temperature of the semiconductor is raised high enough, the free ions can get separated for a short period of time and "hop" from one impurity ion to another; but otherwise they stay fairly near their oppositely charged ions.

I would expect the same is true in solutions, ions stay near each other and only switch when charges (electronegativity) is greater or thermal agitation gives them enough energy to escape the ion they are attracted to.

Note: In the youtube video I gave, I think he said sodium sulfite, and perhaps thats different than sodium metabisulfite. I was thinking that they were the same and metabisulfite has a spare hydrogen atom to create free sulphuric acid. I think that's probably most of my misunderstanding of what you said.

If metabisulfite has two sodiums, then my original thought was misguided; there wouldn't be free hydrogen to from sulfuric acid with in the first place. The stochiometry is good for pure sodium sulfate which is pretty close to neutral. Sodium meta-bisulfate apparently has a pH of around 3.5 to 5. So, it's apparently an acidic salt.

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In your suggestion, the sodium is stripped by the sulfate to make sodium sulfate and copper II chloride. ... I mean, will copper I chloride remain a precipitate with free sulfuric acid around? or will some ppt re-dissolve into solution?

OK. you've said something here very detailed.

The conjugate base for chloride is basically non-extant, for the pka is like < -3; so I follow you about the weak base point.

But if "Chloride" ions stabilize the copper (I), then let me ask a different question;

qualitatively, if I had pure white powder CuCl in the bottom of a flask, no oxygen or air allowed in, and I added HCl to the water; How do I know that the HCl won't be able to oxidize the copper to the +2 state, and release hydrogen gas? What's the theoretical grounds for it reacting or not reacting?

Likewise, what's the theoretical reason that we know that adding bisulfite as free acid, something like: [Na+] [H+] [S2O5]-- won't release the hydrogen to oxidize the copper to the +2 state?

As an abstraction of both the previous ideas, I'd ask the same essential question like this:
If I add excess HCl acid to the solution containing metabilufite, why wouldn't the free acid be able to release hydrogen and attack the CuCl?

With that being said, I know that the equation as described by you will work; I'm just trying to understand if it would work if the pH of the solution was reduced by a free acid's [H+] ions.

[Edited on 30-11-2017 by semiconductive]

DraconicAcid - 30-11-2017 at 00:52

I'm more concerned with safely producing ammonium chloride without making ammonia and chlorine gas... probably something to do outside!

No, just mix dilute hydrochloric acid with dilute aqueous ammonia. NEVER mix chlorine in any form with ammonia- nitrogen trichloride is touchy stuff (Pierre Dulong, who first made it, lost two fingers and an eye to it).

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I don't think it's all that strange you didn't smell ammonia. I made a copper oxide powder (black), and placed it in aqeous ammonia (15%) with oil over it to prevent air from getting in. The ammonia did not dissolve the copper oxide, as far as I could tell, even after a week of sitting. But, after removing the oil the same bottle exposed to air began forming an indigo colored solution within minutes especially if I blew across the water. What is going on is that CO2 from the air attacks the copper oxide and the aqeuous ammonia; producing ammonium carbonate and possibly copper carbonate. The ammonia odor becomes very faint over time, as complexed ammonia with copper is more stable than ammonia carbonate by itself.

Carbon dioxide has nothing to do with this. Copper oxides will dissolve slowly in ammonia to give either the dark blue tetramminecopper(II) ion or the colourless diamminecopper(I) ion (along with hydroxide ion). In your case, it appears there was enough unreacted copper or copper(I) oxide that the latter was formed. Exposing it to air caused oxidation to the copper(II) ion, which is very strongly coloured.

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I would think that the ammonia chloride in your solution would also make a complex with copper chloride, as it forms. So, there ought to be both copper chloride (hydroxide?) held in solution by ammonia, and your precipitated crystals.

Ammonium chloride (not ammonia chloride) can't form a complex- it takes free ammonia molecule to coordinate to a metal (it needs a free lone pair of electrons to act as a ligand).

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I'm pretty disgusted with my pyrex volumetric flasks. Before I do your experiment, I am going to flame polish the bottom and then see if I can deform the concave portion of the bottom to correct the -0.200ml error in volume that I measured. The flask says +-0.05ml, but I'm getting a lot bigger error than marked. If you have any suggestions for how to accurately deform glass ... I'll be thinking about ways to do that tonight.

I suspect that will only make them worse.

DraconicAcid - 30-11-2017 at 01:10

No. When an ionic compound like sodium chloride or sodium sulphate are dissolved in water, the ions become hydrated. The sodium ions hang around with the water molecules, and have very little to do with the chloride ions unless the concentration is very high. I wasn't trying to point out that sodium ions don't have brains- I always anthropomorphize ions and atoms because I'm usually explaining these things to kids. The sulphate ion in a solution of sodium sulphate will act identically to a sulphate ion in a solution of ammonium sulphate, potassium sulphate, cesium sulphate, or even aluminum sulphate- the ions wander far enough away that the sulphate has almost no interaction with the cation.

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Note: In the youtube video I gave, I think he said sodium sulfite, and perhaps thats different than sodium metabisulfite. I was thinking that they were the same and metabisulfite has a spare hydrogen atom to create free sulphuric acid. I think that's probably most of my misunderstanding of what you said.

They are different, but related. Sodium metabisulphite (Na2S2O5) will react with water to give aqueous sodium hydrogen sulphite (NaHSO3)

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Sodium meta-bisulfate apparently has a pH of around 3.5 to 5. So, it's apparently an acidic salt.

Correct

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But if "Chloride" ions stabilize the copper (I), then let me ask a different question;

qualitatively, if I had pure white powder CuCl in the bottom of a flask, no oxygen or air allowed in, and I added HCl to the water; How do I know that the HCl won't be able to oxidize the copper to the +2 state, and release hydrogen gas? What's the theoretical grounds for it reacting or not reacting?

If you look up the reduction potentials of copper, you'll find that they are positive for both copper and copper(I), which means that H+ will not oxidize either of them.

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Likewise, what's the theoretical reason that we know that adding bisulfite as free acid, something like: [Na+] [H+] [S2O5]-- won't release the hydrogen to oxidize the copper to the +2 state?

Because H+ isn't a strong enough oxidizing agent to turn either copper or copper(I) into copper(II).

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If I add excess HCl acid to the solution containing metabilufite, why wouldn't the free acid be able to release hydrogen and attack the CuCl?

Because the acid (H+) won't be able to react with the chloride (it's too weak of a base) or the copper(I) (it's too weak of a reducing agent).

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With that being said, I know that the equation as described by you will work; I'm just trying to understand if it would work if the pH of the solution was reduced by a free acid's [H+] ions.

H+ ions shouldn't make much difference, as long as it's dilute enough that it's still an aqueous solution (and not 90% sulphuric acid).

semiconductive - 30-11-2017 at 09:53

I'm more concerned with safely producing ammonium chloride without making ammonia and chlorine gas... probably something to do outside!

Thanks for all the descriptive input. As I learn more, I'll definitely work my way through all the things you have said here. I was never planning to mix chlorine with ammonia, I'm just aware than mixing aqueous solutions of bleach and ammonia produce both ammonia gas and chlorine gasses, that have hurt lots of house-keepers cleaning toilets. With my limited experience, I wasnt' sure if HCl(aq.) and ammonia (aq.), could also react to release chlorine or ammonia gasses. Case in point, Pierre didn't know what he was doing either ... so better safe than sorry.

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I suspect that will only make them worse.

It's only one flask which seems to be so far "off".
That's why I'm remeasuring with distilled water, and R.O. water after heating to degass them to make sure I didn't make some kind of mistake. If it's not micro gas bubbles which have caused the problem, It could only be a major factory defect, or extreme abuse by a previous owner, that would make the flask be off by this much.
I checked the calibration on the mass scale, and it's fine. I'm only 100 feet above sea level, so I assume that I'm close enough to standard pressure to not affect 50cc's of waters volume by 0.2cc.

If I haven't made a mistake, then my only choices are to re-mark the neck with a second precision scribe (hard to do), or soften the glass on the bottom of the flask enough to make it less concave. Ordering another flask is always possible if I totally ruin it, but that takes a lot more time and money. So, I'm willing to give fixing it a shot as I have nothing to loose.

[Edited on 30-11-2017 by semiconductive]

DraconicAcid - 30-11-2017 at 11:34

I think your best best would be to re-mark the volume with a very thin piece of electrical tape or duct tape. It's easy, close enough, and not going to come off under normal washing conditions.

DraconicAcid - 30-11-2017 at 12:09

I was never planning to mix chlorine with ammonia, I'm just aware than mixing aqueous solutions of bleach and ammonia produce both ammonia gas and chlorine gasses, that have hurt lots of house-keepers cleaning toilets. With my limited experience, I wasnt' sure if HCl(aq.) and ammonia (aq.), could also react to release chlorine or ammonia gasses.

That's a myth. Acids mixed with bleach will produce chlorine gas; ammonia mixed with bleach will produce chloramines, which are also toxic.

semiconductive - 30-11-2017 at 14:02

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If you look up the reduction potentials of copper, you'll find that they are positive for both copper and copper(I), which means that H+ will not oxidize either of them.

In the table of reduction potentials, copper solid metal to Cu++ is +0.34V, and Cu+ is +0.52V.

I also notice that more electric field / potential energy is required to cause a reduction of the single ionized copper, than that of the doubly ionized copper atom.

That's odd ... intuitively, I'm used to working with atomic calculations in vacuum. removing the first electron generally takes less electric field/energy than removing the second electron.

In the reduction potential chart, I notice that Ag+'s reduction is at 0.8V, and Ag++'s at 1.98; which follows the same trend as I am used to in vacuum. That suggests to me, that water may have an "average" reference voltage that is analogous to a vacuum/ground reference at infinity for normal physics and electronics calculation purposes.

I am guessing it's going to be be a potential > about +0.74 volts with respect to a hydrogen electrode as that's the minimum voltage required to cause all ions on the reduction potentials charts (that I can see) to agree with the principle that first ionizations are always lower magnitude to the reference voltage than the second ions.

In any event, copper is easier to oxidize to the +2 state, than it is to the +1 state with respect to a hydrogen electrode and that's counter-intuitive; but I can accept it. Reference voltages are arbitrary.

However, that suggests copper will not be attacked at all by hydrochloric acid.

When I take copper, and put it in hydrochloric acid; it will generally clean off the oxide very quickly, if any exists. But even after the copper is clean, hydrochloric acid will continue to attack copper slowly .. and faster as the temperature is raised.

The prediction of the reduction potential chart seems hard to believe/wrap my head around. I use HCl to etch copper circuit boards.
So, what you're telling me is something I'm going to have to test out.

I think I have to place the copper in HCl with non access to oxygen, air, to prevent oxidation (silicone oil or freon capping the liquid in an air-tight way) and according to the chart that would suggest HCl could no longer attack the copper metal at all at room temperature (21C) at my house.

This also brings up the problem of temperature that I've been wondering about.
In semiconductor physics, the law of mass action can be derived by taking the fermi level into account and the density of states of the secmiconductor in "band" theory. I forget exactly how it was done, but the ions concentrations being simple products was an approximate result of the derivation; but it was in fact dependent both on temperature and ion concentration (slightly.)

In Chemistry, I think the law of mass action is just accept a-priori / empirically ... but the chemical potentials are going to have to change significantly with respect to temperature; eg: something like:

~ - k1 * RT ln ( k2 )

I know batteries change by millivolts/degree Celsius; which means the oxidation/reduction potentials must also vary with temperature. I note the chart in my book says at "25C", so that's suggestive that the chart is increasingly invalid at other temperatures.

Just as a crude low-balled order of magnitude (guess) of 4millivolts/*C would mean that the Cu +0.34V potential could be totally overcome by a change of 85 degrees Celsius. If the temperature change is in the right direction. Seeing that HCl etches copper faster as it gets hotter, I would expect that heat could cause the lowering of the threshold that the chemical reaction takes place at.

[Edited on 30-11-2017 by semiconductive]

DraconicAcid - 30-11-2017 at 15:17

In vacuum, this is true, but in solution, the water more strongly coordinates the divalent cation and stabilizes it more. Copper(I) ions really don't exist in aqueous solution, unless they are coordinated to a ligand that stabilizes them more than water does. Otherwise, 2 Cu(+) -> Cu(2+) + Cu(s).

There are many metals which do not form stable +1 ions, and thus the ions with higher charges are easier to make. You will never find a reduction potential for, say, zinc(I) or chromium(I)- these ions don't exist.

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That suggests to me, that water may have an "average" reference voltage that is analogous to a vacuum/ground reference at infinity for normal physics and electronics calculation purposes.

I do not know what you mean by this- water is not analogous to a vacuum in any way.

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I am guessing it's going to be be a potential > about +0.74 volts with respect to a hydrogen electrode as that's the minimum voltage required to cause all ions on the reduction potentials charts (that I can see) to agree with the principle that first ionizations are always lower magnitude to the reference voltage than the second ions.

Aqueous ions won't agree with that principle, otherwise you would find aluminum(I) and aluminum(II) on your tables of reduction potentials.

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However, that suggests copper will not be attacked at all by hydrochloric acid.

When I take copper, and put it in hydrochloric acid; it will generally clean off the oxide very quickly, if any exists. But even after the copper is clean, hydrochloric acid will continue to attack copper slowly .. and faster as the temperature is raised.

The hydrochloric acid will clean off the oxide very quickly- this is an acid-base reaction and is not affected by the redox potentials.

Once the copper is clean, you have copper ions and chloride ions in solution. In the presence of a high concentration of chloride ions, copper(I) forms a stable complex ion called dichlorocuprate(I) [CuCl2]-. This will react with oxygen in solution or in the air to form copper(II) ions, which then react with metallic copper in the presence of chloride ions to give more dichlorocuprate(I) ions....which then react with oxygen.

It is not the reaction Cu + 2 HCl -> H2 + CuCl2, which is thermodynamically unfavourable (as indicated by the redox table).

It is the overall reaction 2 Cu + 4 HCl + O2 -> 2 CuCl2 + 2 H2O, which is autocatalyzed by the presence of CuCl2.

This reaction will happen faster at higher temperatures because all reactions happen at higher temperatures, unless it's a process disfavoured by high temperatures (like the freezing of water).

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In Chemistry, I think the law of mass action is just accept a-priori / empirically ... but the chemical potentials are going to have to change significantly with respect to temperature;

We can derive the law of mass action from thermodynamics.

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I know batteries change by millivolts/degree Celsius; which means the oxidation/reduction potentials must also vary with temperature. I note the chart in my book says at "25C", so that's suggestive that the chart is increasingly invalid at other temperatures.

That's true, they will change with temperature. If you have the reaction enthalpy, you can calculate the equilibrium constant at any other temperature easily enough, and from there you can find the reaction potential.

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Just as a crude low-balled order of magnitude (guess) of 4millivolts/*C would mean that the Cu +0.34V potential could be totally overcome by a change of 85 degrees Celsius. If the temperature change is in the right direction. Seeing that HCl etches copper faster as it gets hotter, I would expect that heat could cause the lowering of the threshold that the chemical reaction takes place at.

I'll do the calculation for you when I have time, but i doubt it will work that way. It's just a kinetic effect of reactions happening faster at elevated temperatures.

semiconductive - 14-12-2017 at 20:31

Quote: Originally posted by Chemetix

Copper (II) sulphate + NaCl => Cu(II)Cl2
Edit:
I know you wanted to use reducing sugars but with such strong ions like Cl- sugar isn't going to be able to pull that off, tollens reaction uses a fairly weak cation like silver and ammonia to basically balance the force of the nitrate then the sugar/aldehyde can get in there and do it's thing.
[Edited on 29-11-2017 by Chemetix]

@Chemetix,
I Missed this comment from before.

How do I find out the strength of a reducing sugar, compared to Cl- , OH-, and HCO3- ?
More importantly, why would it need to remove Cl-, when the copper atoms are not oxidized with two chlorines in the first place, but only one?

I only added enough HCl to attack half the basic copper carbonate bonds;
Did I miscalculate the molar concentration of HCl? Would you check for me?

Basic copper carbonate is in the Cu+2 oxidation state. The +2 is caused by BOTH hydroxil anion, and a carbonate anion. I didn't supply enough HCl to replace all the CO2 and OH. Therefore, one of the bonds must still be attached to the copper ... because the copper did all dissolve into solution once the acid was added; I don't think I created a solution where chloride attacked only some of the basic copper carbonate, and left part of it alone. Otherwise, why did all the copper carbonate go into solution?

I assumed either the hydroxide or the carbonate would be easier to displace, and the copper atom would stay in the +2 oxidation state because one of the anions was HCl, but the OTHER anion would still be either hydroxide or carbonate.

My hope was that a reducing sugar would have an advantage in a situation where chloride was NOT the only anion attached to the copper, but a weaker anion was present.

Besides, what you've told me is inconsistent or grossly over-simplified compared to what other people are saying.

The Wikipedia article I was following to make copper I chloride explicitly said that reducing sugars (eg: ones LIKE ascorbic acid / vitamin C ) was able to reduce copper II chloride to copper I chloride. So, they are saying a reducing sugar can even overcome chloride. But you say it can't?

Wikipedia doesn't specify under what conditions it happens. (eg: Do you have to dry out the solution, or does it precipitate, etc. )

The article only said on December 14,2017

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https://en.wikipedia.org/wiki/Copper(I)_chloride
"Copper(I) chloride can also be prepared by reducing copper(II) chloride, e.g. with sulfur dioxide or a reducing sugar such as Ascorbic Acid (Vitamin C):

2 CuCl2 + SO2 + 2 H2O → 2 CuCl + H2SO4 + 2 HCl
2 CuCl2 + C6H8O6 →2 CuCl + 2HCl + C6H6O6

Many other reducing agents can be used.[14]"

My thought is that was that ascorbic acid is a reducing sugar because it has the ability to change oxidation states from dehydroascorbic acid <--> ascorbic acid. It's a simple hydrogen/carbon/oxygen chemical. Glucose is very similar. From what I can see from wikipedia, The important part of ascorbic acid is the hydroxil groups on the ring structure of the reducing sugar which participates in oxidation and reduction. They aren't even part of an aldehyde group.

Glucose can form a ring structure with no aldehyde, or it can temporarily open and present a very active aldehyde group. So, Glucose has both the necessary hydroxil groups to act the same as ascorbic acid (AKA: a reducing sugar.) and the ability to act as an aldehyde. So its a hexose, and ascorbic acid looks like a hexose as well; they both have 6 carbons, and 6 oxygens, They both even form a ring structure.

It's not like I can look the oxidation and reduction potentials of reducing sugars up in my oxidation/reduction tables. My college texts aren't that comprehensive. As useful as wikipedia is ... it doesn't list a half cell potential or anything like that which I can use to compare (at least yet, December 14th, 2017).

https://en.wikipedia.org/wiki/Vitamin_C
https://en.wikipedia.org/wiki/Glucose

How can I know which reducing sugars will work, and under what conditions?
What's the key point about these reducing sugars that is different? What pattern of atoms do I look for to find a reducing sugar "like" ascorbic acid?

[Edited on 15-12-2017 by semiconductive]

DraconicAcid - 14-12-2017 at 22:03

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How do I find out the strength of a reducing sugar, compared to Cl- , OH-, and HCO3- ?
More importantly, why would it need to remove Cl-, when the copper atoms are not oxidized with two chlorines in the first place, but only one?

You don't need to remove chloride. You can't compare the strength of a reducing agent to that of chloride, hydroxide and bicarbonate, since these are not reducing agents. I have no idea what Chemetix meant by chloride being a strong anion- you're not pulling off a chloride; you're adding an electron.

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I only added enough HCl to attack half the basic copper carbonate bonds; Did I miscalculate the molar concentration of HCl? Would you check for me?

Basic copper carbonate is in the Cu+2 oxidation state. The +2 is caused by BOTH hydroxil anion, and a carbonate anion. I didn't supply enough HCl to replace all the CO2 and OH. Therefore, one of the bonds must still be attached to the copper ... because the copper did all dissolve into solution once the acid was added; I don't think I created a solution where chloride attacked only some of the basic copper carbonate, and left part of it alone. Otherwise, why did all the copper carbonate go into solution?

If it all went into solution, then you neutralized all of the hydroxide, and probably converted a significant amount of the carbonate into bicarbonate. But it's not the chloride that's attacking the basic copper carbonate; the chloride is a spectator ion. The hydrogen ions from the acid are attacking the hydroxide and the carbonate.

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I assumed either the hydroxide or the carbonate would be easier to displace, and the copper atom would stay in the +2 oxidation state because one of the anions was HCl, but the OTHER anion would still be either hydroxide or carbonate.

An acid-base reaction will not change the oxidation state of copper(II).

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My hope was that a reducing sugar would have an advantage in a situation where chloride was NOT the only anion attached to the copper, but a weaker anion was present.

The chloride is *not* attached to the copper in dilute aqueous solution. It's just a counterion.

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The Wikipedia article I was following to make copper I chloride explicitly said that reducing sugars (eg: ones LIKE ascorbic acid / vitamin C ) was able to reduce copper II chloride to copper I chloride. So, they are saying a reducing sugar can even overcome chloride. But you say it can't?

Wikipedia doesn't specify under what conditions it happens. (eg: Do you have to dry out the solution, or does it precipitate, etc. )

CuCl isn't very soluble, so if the reaction occurs, it will precipitate out easily. Try it- if it doesn't go, heat it. Make sure the chloride ion concentration is moderate (i.e., enough to make sure that your product will precipitate out, but not so high that you form soluble complex ions. Aim for about 0.5 mol/L max).

[Edited on 15-12-2017 by DraconicAcid]

semiconductive - 16-12-2017 at 02:44

Quote:

DraconicAcid, your way of thinking is very narrow and restricted only to the solution.

For example, if I took table salt and dissolved it in water ... it's still called "salt water."
I know you are correct ... the chloride anions are not attached to the sodium when dissolved; however, when the water is evaporated, or the salt is precipitated (by adding alcohol, or other chemical to physically change the solubility); the chloride ions would re-attach to the sodium and form "salt" again. The chloride ions are not stripped (or displaced) from the sodium by alcohol or evaporation during the formation of crystals.

Two salts can also be added to water. For example sodium chloride, and copper acetate.

During the evaporation of water, two salts will form again. The particular salts that will form will be in proportion to the cation and anion strength. I've done experiments like this, and unless the reaction is a "metastatis" reaction, the original two salts will precipitate out. In general, if the evaporation is slow, distinct and tiny salt crystals will form at random locations in the flask; but each tiny crystal will be one or the other kind of salt and relatively pure. The anions and cations in solution, during evaporation, do have a preference as to what chemical they will bind with during solidification.

When I talked about "stripping" in earlier posts, I was thinking of the end result of the salt formed after evaporation or precipitation. I'm not thinking about what is in solution.

People have a tendency to think about the end product. It's possible that's what Chemtix was thinking as well. It's not precise, but it's not good to get hung up over that kind of language.

Quote:

I didn't add enough HCl to neutralize all the hydroxide and carbonate.
Basic copper carbonate is: OH-Cu-CO3-Cu-OH
(note: It won't be spectator ion once evaporation or precipitation happens.)

The stoichometery of the first post was intended to add exactly the number of moles of HCl that I had moles of copper. Therefore, since the basic copper carbonate has one hydroxide each, and 1/2 carbonate each, there isn't enough Hydrogen to make bicarbonate AND neutralize the hydroxide ions.

So, would you check the math of the first post. Because, your conjecture is impossible unless I made a gross math mistake.

The original chemical reaction mixture has been stirring at 55C for over two weeks. A very small amount of light orange colored precipitate did form at the top surface of the sealed Erlenmeyer flask. No more oxygen can get in, but the orange precipitate has sunk to the bottom after about a week; It's immeasurably small and not even a thick enough film to cover all the glass on the bottom of the flask. (Edit: This remark is slightly incorrect, see photos in posts which immediately follow this one. The orange film is not enough to make the glass opaque.) The original carbonate was over two millimeters thick covering the flask bottom before being dissolved.

If the carbonate is not displaced during evaporation/precipitaiton, (Eg: assuming the hydroxil is removed/neutralized more easily than carbonate.) then I would expect: Cl-Cu-CO3-Cu-Cl to be the product of evaporation.

Quote:

An acid-base reaction will not change the oxidation state of copper(II).

Right, which is exactly what I just said.

Quote:

My hope was that a reducing sugar would have an advantage in a situation where chloride was NOT the only anion attached to the copper, but a weaker anion was present.

The chloride is *not* attached to the copper in dilute aqueous solution. It's just a counterion.

But it will be attached when I remove the water; eg: if I put a vacuum pump on it and evaporate the water without allowing oxygen into the solution. Note, I did say that wikipedia didn't tell me what conditions the reducing sugar had an effect in.

Quote:

The glucose solution has already been baking for 12 days at 55*C. That temperature was chosen because another wikipedia article talked about brine solutions changing copper I into a different compound at 60*C +. So, I chose a slightly lower temperature to not trigger a side reaction.

The number of moles of carbonate is ~0.008 dissolved in 50mL of water. So, double that number of moles of copper is present inside the carbonate: ~0.0016 mole /0.050L = 0.032 which is << 0.5 mol/L. So, your requirement is already met.

I could put ascorbic acid into another beaker, and I intend to do that. But doing the experiment isn't going to explain to me *WHY* ascorbic acid ( a reducing sugar ) works when glucose (another reducing sugar) doesn't.

eg: the original question about strength of the reducing sugar vs. carbonate or chloride could be rephrased: How do I compare the strength of the two reducing sugars ?

I really don't see how it helps to think about what's going on in solution when I can't really test it. At some point, for reduction to have happened; chloride ion is prevented from attacking the copper during evaporation because of the presence of an electron transfer, or a hydrogen atom has reduced copper to the +1 state.

What matters is not so much "when" the reduction occurs; rather what I'm focused on is that Wikipedia says a reducing sugar can cause Copper I to be made from copper II . Wikipedia was talking about a solution of CuCl2. In order for that to be reduced, at some point .. one of the chlorides had to find something OTHER than copper to be attracted to during evaporation of water or during precipitation. (eg: Chloride finds hydrogen and turns into a gas and evaporates.)

EDIT: In the footnotes of wikipedia, they cited an old text book which happens to be hosted by sciencemaddnes.org, but when I tried to access the book, the PDF reader comes up, but the document never shows. Hence, I can't read the conditions of the original experiment with ascorbic acid.

[14] Glemser, O. and Sauer, H. (1963) "Copper(I) Chloride" in Handbook of Preparative Inorganic Chemistry, 2nd ed. G. Brauer (ed), Academic Press, NY. Vol. 1. p. 1005.

http://www.sciencemadness.org/library/books/brauer_ocr.pdf

I would appreciate someone checking the original book for me and telling me the conditions of the experiment that wikipedia is citing.

[Edited on 17-12-2017 by semiconductive]

semiconductive - 16-12-2017 at 16:16

Quote: Originally posted by Boffis

If you have time the simplest method is to place bare copper wire or bits of copper in a jar and fill it up with dilute ammonium chloride solution and seal tightly.

I bought the two dram vials from e-bay (~50cents/piece). I thought I had purchased the borosilicate glass ones, but now that I check the listing, it doesn't specify the kind of glass; (seller: homeluxe-grandparfumes ,8mL 1/4Oz Empty Glass Bottle Screw Top Clear Sample Vial Perfume Oil 2 Dram ) So I bought the wrong listing, accidentally. I'll heat one later today to check the coefficient of expansion of the glass and see if it's lime glass or borosilicate; but I'm sure it has a phenolic cap. Phenolic caps can handle up to 220*C, so it's a pretty decent vial for doing experiments in. I'm curious about what kind of container you did the experiment in.

I measured my vials dimensions at 21.5C for the glass (not the cap). There is a photo of that + an ABS pipe I drilled to hold 12 vials with the insides of two of the caps exposed:

2dram.png - 5kB

It should have more than 8mL of storage, eg: around 11 to 14, depending on glass thickness. I can't quite measure the inside. I will probably measure the volume by weighing water later today.

As you can clearly see, some of the phenolic caps come with a foam sealing liner, but some of them don't. The manufacturer is kind of lazy and inconsistent.

When you did your experiment, what kind of container did you seal the chemicals in; and was it an air-tight seal with some way to exclude air bubbles? eg: Did you do anything to make sure the chemicals were in an inert environment? eg: teflon, etc? and, especially, Do you have any idea of how much wire and ammonium chloride that you used?

Also, was the container exposed to light; or was it kept in a dark, even temperature place (crystal growth is often affected by light or thermal gradients. )

[Edited on 17-12-2017 by semiconductive]

Original experiment photo update.

semiconductive - 16-12-2017 at 16:42

It's been two weeks. Here is a photo shoot showing the Dec 16,2017 state of the glucose and copper carbonate/HCl solution described in the first post of this thread. The first picture, is a copy of the last picture in the opening post of the thread so you can compare colors. All pictures are shot with the same cell phone camera. warning: To my eyes, the actual color of the solutions was slightly more green than the computer or cellphone screen shows. The CCD camera in the phone tends to emphasize blue over green; still,all photos are with the same phone, so the differences in shades between photos is accurate relative to the other photo.

one.jpg - 16kB

As you can see, the solution color has turned much greener over the last 14 days. A very fine orange colored film is adhering to the glass of the flask. At first, I thought it was just on the surface of the solution, and had sunk to the bottom of the flask. But now, since I took the flask off the magnetic stir plate, I can clearly see the orange film adhering to the glass of the flask and being reflected by the glass on the bottom of the flask. That film could easily be either copper metal, or something like copper I oxide. Cu-O-Cu. It's hard to tell. There was a tiny bit of the film floating on top of the solution several days ago, but it's no longer there. The film appears to be entirely coating the glass.

It's important to note that the air head-space of the Erlenmeyer flask is sealed under a glass plug. The plug has not been removed at any time in the last two weeks, so only the air in the head-space could have supplied any oxygen, and it is very limited in quantity. I also have reducing sugar in the solution, so hopefully the dissolved oxygen is not really causing oxidation.

The only other unusual issue is that the teflon stir bar (not clearly visible in pictures) has a black spot on the part that rubs against the glass on the bottom of the flask. When I shook it and got the stir bar to flip over, the black mark on top of the bar dissolved while another black mark formed on the bottom of the bar. I have no idea what would cause that.

I set the constant temperature at 55*C, and covered the flask with refractory since the experiment began. The original refractory brick chunks were not very effective in keeping the entire flask at the same temperature as the hot plate. I can't put a thermometer in the flask without introducing a stainless steel rod, or leaking air. So, the early part of the experiment's liquid had an average temperature below 55*C. The temperature probe is actually measuring the hot plate outside the flask, just next to the flask.

I'm lifting a new insulation jacket in the photo. It's made of cheap (disposable) plaster of Paris and pearlite. I will talk about improving that more in another thread, as my first casting isn't very professional. In other experiments I've had it weaken, crumble and crack when heated above 150C in a kitchen oven. However, when kiln heat like like 700C to 900C is applied, plaster of Paris can become really hard and tough after shrinking. It's fickle stuff to make molds and shaped objects from that can take heat consistently ... I'd like to fix that. At least, plaster/pearlite is a much better insulator than aluminum foil which is typically shown by home chemists on youtube.com to wrap flasks during heating for distillation.

[Edited on 17-12-2017 by semiconductive]

zed - 18-12-2017 at 15:07

Ummm. I'm not very careful, but when I mix CuSO4 solution and Ascorbic Acid, and return a week or so later, I find nano particles of copper, and a clear solution. That's it. Now, if there is Cu 1+ remaining in the solution, I haven't checked.

Got experimental details, from some folks that claim to have the Cu1+ thing..... pretty well wired.

https://www.dimanregional.org/site/handlers/filedownload.ash...

[Edited on 18-12-2017 by zed]

DraconicAcid - 18-12-2017 at 15:17

Quote: Originally posted by zed

There won't be. You need an anion that forms an insoluble compound with copper(I) to stabilize it- chloride or bromide, for example.

DraconicAcid - 19-12-2017 at 12:41

Have you checked the pH of that solution?

Anyway, I found a table that gives reduction potentials for some organic things:

Dehydroascorbic acid + 2 H+ + 2e -> ascorbic acid E = + 0.390 V.

This tells me that the overall potential for AA + Cu(2+) --> DHA + 2 H(+) + Cu(s) is -0.053 V, so thermodynamically unfavourable under standard conditions. If you reduce the concentration of hydrogen ions, that will make the reaction more favourable, and it will be spontaneous at a pH of 2 or greater.

The table doesn't have any sugars on it, but it does have HCO2H + 2H(+) + 2e --> H2CO + H2O E = +0.056 V, so if we declare that all aldehydes will be basically the same, we can say that:

RCHO + H2O + Cu(2+) --> RCO2H + Cu(s) + 2 H(+) will have E = 0.28 V and thus be spontaneous (and the lower the concentration of H+, the more spontaneous it should be).

(As an aside, I did once heat copper sulphate in a concentrated solution of sugar, thinking the higher viscosity might favour larger crystals. It gave copper powder.)

If you have chloride in solution, you should have no trouble getting CuCl forming instead of copper metal.

ETA: I dissolved 0.15 g CuO in 2 mL of 6 M HCl, added 1 mL of 1 M NaOH and 0.7 g of glucose in distilled water, and have been heating in a hot water bath. I will report any changes.

ETA: Nothing happened for a while, so I added 1.6 g sodium chloride and adjusted the pH to about 6 with acetic acid and sodium hydroxide- lots of greenish precipitate.

ETA: Said precipitate dissolves readily in dilute HCl, and therefore must have been either the hydroxide, or basic chloride. I'll try again tomorrow, adjusting the pH to keep the copper in solution first, then adding glucose.

[Edited on 20-12-2017 by DraconicAcid]

semiconductive - 21-12-2017 at 16:27

I broke the seal on the flask, and took a sample of about 1CC, and used about 1/10th a soil pH test pill meant for 10 CC of water. The solution color changed to slight purple, which is not an indicator color ... but since the color was blue-green to begin with, the color of the pH indicator can not be adding green but must be adding tan/orange/red to the solution.

see photo:
From top of the color chart to bottom pH is: 7.5, 7.0, 6.5, 6.0, 5.5, 5.0, 4.5 (or below).
The pill powder is white. Looking at the powder near the bottom of the container shows a tan discoloration, not quite orange. That puts the pH at about 5.5 to 6.

Edit: ( I diss a suppress blue channel in top of the liquid, to give a better idea of color since blue is not part of pH scale. I would estimate a pH of 6 to 6.5 with modified color. )

At my house, when reverse osmosis or distilled water is exposed to air for 24 hours I get about a pH of around 5.5. The pH of the copper solution is pretty much in the same range, so I assume that actual solution is either neutral or slightly acidic depending on how much CO2 is dissolved in the water.

Copper sulfate precipitated with sodium carbonate normally doesn't form pure carbonate, but from what I'm told it forms copper hydroxide carbonate. OH-Cu-CO3-Cu-OH.

But, I don't see any reason that copper couldn't also form longer chains with carbonate...
eg: something like: OH-Cu-CO3-Cu-CO3-Cu-OH

The longer the chain, the higher the density of the molecule; and the more likely it would be to settle to the bottom of the flask or precipitate.

Hydrogen ions, whether they are from HCl or from carbonic acid will act the same in solution regardless of where they came from; but for various salts and acids, there are different solubility and ionization products determining dynamic equilibrium.

Eg: some acids, like sulfuric, have an ionization constant that is very high for the first proton (hydrogen atom) removed. But a much smaller ionization product (if I recall correctly) for the second proton. So, sulphuric acid would have 100% 1st ionization, but not quite 100% second ionization. H-Cl is pretty much 100% ionized all the time. That's the nature of strong acids.

However, carbonic acid is a weak acid and shouldn't we treat it differently / theoretically?

Why is is that copper forms basic copper carbonate, rather than pure copper cabonate?
Would it form pure copper carbonate in solution if I added pressuized CO2 to the liquid?
( Calcium carbonate, for example, can be dissolved in carbonic acid ).

eg: Isn't it likely that some acids are less likely to be ionized by water and separated from the copper atoms than others?

In the experiment I did, is it possible that the hydroxides of basic copper carbonate have been neutralized into water from the hydrogen supplied by HCl, but the remaining carbonates are not disassociated from the copper at all?

Is there such a thing as copper (I) carbonate? Cu-CO3-Cu
Or is copper always ionized in water to +2, so that H+ OH- so that the solution is:
+Cu-CO3-Cu+

What's the difference between adding HCl, to remove the OH radiacls by supplying free H+ ions, and supplying an electron to reduce copper to the +1 oxidation state?

How do we know which ion species are soluble in water, and which aren't?
eg: the hydroxide carbonate (OH-Cu-CO3-Cu-OH) is not soluble; But,, does that mean that when I add enough HCl to make: Cl-Cu-CO3-Cu-Cl (on drying); that this chloride half salt is insoluble? ( I am assuming hydroxide is more easily displaced than carbonate. )

Edit: The pH indicator powder color is getting closer to orange with time, so that means the solution exposed to air is becoming slightly more acidic with time.
(Picture after several hours, I electronically removed the offending color blue .. since the chart is purely red-green, should make it easier to tell ph. eg: Between 5.5 and 6. )

[Edited on 22-12-2017 by semiconductive]

DraconicAcid - 21-12-2017 at 18:22

Copper sulfate precipitated with sodium carbonate normally doesn't form pure carbonate, but from what I'm told it forms copper hydroxide carbonate. OH-Cu-CO3-Cu-OH.

But, I don't see any reason that copper couldn't also form longer chains with carbonate...
eg: something like: OH-Cu-CO3-Cu-CO3-Cu-OH

The longer the chain, the higher the density of the molecule; and the more likely it would be to settle to the bottom of the flask or precipitate.

This is an ionic compound, not a polymer, and not a molecule. You have a regular arrangement of copper ions, carbonate ions, and hydroxide ions that repeats over and over and over again. The unit cell looks like this:

Now imagine that as a brick, and you've got a vast number of these bricks stacked together into a block that's billions of blocks long, billions of blocks wide, and billions of blocks high. That's a very tiny crystal of basic copper(II) carbonate. The bigger the crystal, the better it will settle, but that has more to do with mass and solution viscosity than density.

Quote:

Hydrogen ions, whether they are from HCl or from carbonic acid will act the same in solution regardless of where they came from; but for various salts and acids, there are different solubility and ionization products determining dynamic equilibrium.

Eg: some acids, like sulfuric, have an ionization constant that is very high for the first proton (hydrogen atom) removed. But a much smaller ionization product (if I recall correctly) for the second proton. So, sulphuric acid would have 100% 1st ionization, but not quite 100% second ionization. H-Cl is pretty much 100% ionized all the time. That's the nature of strong acids.

That is correct.

Quote:

However, carbonic acid is a weak acid and shouldn't we treat it differently / theoretically?

Why is is that copper forms basic copper carbonate, rather than pure copper carbonate?
Would it form pure copper carbonate in solution if I added pressuized CO2 to the liquid?
( Calcium carbonate, for example, can be dissolved in carbonic acid ).

eg: Isn't it likely that some acids are less likely to be ionized by water and separated from the copper atoms than others?

I'm not actually sure why copper forms a basic carbonate rather than a simple one; it also tends to form a basic chloride and a basic sulphate rather than a simple precipitate of the hydroxide. I've read that you can crystallize a non-basic copper(II) carbonate from a saturated solution of sodium carbonate, but I've never tried it.

There are some anions which will coordinate to copper ion solution, particularly if the anion is concentrated. You can make ions such as CuCl4(2-). Carbonate is a lousy ligand, but will coordinate to copper in an extremely concentrated solution.

Quote:

In the experiment I did, is it possible that the hydroxides of basic copper carbonate have been neutralized into water from the hydrogen supplied by HCl, but the remaining carbonates are not disassociated from the copper at all?

No- as I said, carbonate is a lousy ligand.

semiconductive - 21-12-2017 at 22:45

I'm preparing to repeat the experiment. My intention is to use ascorbic acid as reducing sugar, and carefully test the dissolution of basic copper carbonate in HCl to insure I don't have excess HCl.

I think I made a math mistake in the first post, though it doesn't change the conclusions.
The number of moles I computed was slightly off.
I'd appreciate someone double checking my molecular weights and formula.

I think the following is the reaction to produce basic copper carbonate:

2CuSO4 + 1H2O + 2Na2CO3 --> 2Na2SO4 + Cu2CO3[OH]2 + CO2

This time I weighed 2.508g of CuSO4 * 5H2O @ 249.685g/mol --> 0.01004 mol
I added Na2CO3, (Washing soda, not bicarb) until the water went almost perfectly clear,
A large amount of bubbles were formed, so it's not a simple metastasis reaction; but hydroxide is being absorbed from the water and gasses are being released (CO2)

I added filter paper to a Buchner funnel, weighing 0.218mg. I was hoping to be able to weigh the filter paper and get a more accurate total conversion mass. (To track losses of chemical). Then I vacuum filtered the solution with a single rinse of 50ml water. I won't do this again in the future as getting all the carbonate out was difficult.

I tried heating the Buchner funnel and paper afterward to dry the carbonate, but I used too much heat and some black oxide formed in the carbonate, making it impure; so I'll have to re-do the experiment tomorrow to get a precision measurement.

Basic copper carbonate is supposed to have a mass of 221.113g/mol ; and two moles of copper sulfate, should produce one mole of basic copper carbonate. (Therefore, 0.01004mol --> 0.00502 mol) I should have a maximum of 1.1105 grams of precipitate. However, I got almost 1.4grams after subtracting the filter paper weight, which is way to high ( 0.0057 mol. ) The filtrate wasn't perfectly dry, so it could also be moisture.

In any event, I was mostly interested in how much acid it would take to dissolve all the copper carbonate (0.005 moles worth.), so I weighed out 1.904 grams of HCL solution at 31.45% concentration which should be ~~0.599g HCl @ 36.461g/mol, --> 0.0164 moles. That's the same amount I used last time, and can neutralilze the hydroxide (but not all the carbonate) of 0.0082 moles of basic copper carbonate. (eg: 0.05 moles of basic copper carbonate). The amount of HCl used, I think, was a mistake in the first post of this thread.

However, this time when I added the acid ... I did it without rapid stirring. I wanted to see exactly when all the precipitate was dissolved, so I could weigh the remaining acid. I also wanted to see how much gas (if any) would be released which is hard to see when stirring vigourously. If the acid was only neutralizing hydroxides, there would be no bubbles ... but if it was displacing the CO2, then there would be bubbles.

I added diluted HCl drop by drop (I added water to make 15g total from the 1.9g I measured ). There was a very small amount of bubbling as I added diluted acid, but no where near the volume of CO2 released when making the basic CuCO3[OH2] from washing soda. I noticed that the black oxide dissolved first, and the carbonate dissolved last. I expected to stop adding acid when all the carbonate dissolved, but even with all the HCl added, there is still a small amount of carbonate left on the bottom of the flask.

That's when I realized that copper oxide takes more acid to dissolve than does carbonate. CuO --> CuCl2. Since some carbonate was escaping due to not stirring, I was producing more CuCl2 than the last time I did the experiment.

So, I'll try again tomorrow and have someone double check all my measurements and steps. How quickly acid is added, and the amount of copper oxide present, and stirring rate obviously affects the amount of carbonate that stays in solution.

Note: The HCl is sold as 31.45% solution, but half has been used up and since a little gas fumes out each time I open up the cap ... it's likely lost some HCl strength.

Edit: Adding pictures; Although the quantites and sources of reagents started with (1 gallon muriatic acid jug, root killer CuSO4.5H2O) are the same; the processing was different. The primary differences are that the original experiment added chloride to carbonate at room temperature (21.5C) under vigorous stirring; while the second experiment did it at 55C with little to no stirring at first. The carbonate of the second experiment was also accidentally burnt, which caused some of the carbonate to turn in to copper II oxide. (black.) Sugar will not be added to the second experiment, because it's already defective.

Note: In the first picture, the tare weighted flask is just holding diluted HCl while the stove flask is the second experiment (Tare weight is on back side of flask); but the truly un-marked flask is only on the left in the final two pictures. The unmarked left flask is the original experiment, uncapped, to allow air in at room temp.

The final two pictures are after allowing both experiments to sit over-night in open air: No extra precipitate formed in the original experiment due to air; and the remaining precipitate in the second experiment did not dissolve due to air. Note the color of the precipitate in the right flask looks different when viewed from above (though solution) or below (through glass).

[Edited on 22-12-2017 by semiconductive]

semiconductive - 21-12-2017 at 23:42

I'm looking at it ... but that diagram does not match the molecular formula I have been told represents basic copper carbonate. It also depicts bonds that defy normal oxidation states.

The black atom has to be carbon, and the red ones are oxygen. Therefore, the lighter orange ones are copper and the grey/white are hydrogen.

The molecular formula I learned for basic copper carbonate is: HO-Cu-CO3-Cu-OH
The carbonate CO3, has one double bonded oxygen, and two single bonded oxygens.

Cu-OH
|
O
|
C=O
|
O
|
Cu-OH

Empirical:
1C:2H:5O:2Cu

But, instead of having one copper atom attached to the carbonate on each side, the drawing has two copper atoms. It doesn't make sense to me.
Three bonds puts oxygen in the -3 state .... !?

It's cool looking, but I have no idea what kind of bonds it's trying to depict. Oxygen doesn't make three single bonds.

I can see the way it's supposed to repeat in 3D; and counting the atoms that have bonds:
4C : 8H : 20O : 8Cu Since everything's a multiple of 4, that's 1C : 2H : 5O : 2Cu

The stoiciometry is correct ... however, it's impossible to understand why some of the oxygens have three bonds to carbon/copper or single bonds to hydrogen.
Hydrogen monoxide ?

[Edited on 22-12-2017 by semiconductive]

semiconductive - 27-12-2017 at 17:48

Quote: Originally posted by zed

Ummm. I'm not very careful, but when I mix CuSO4 solution and Ascorbic Acid, and return a week or so later, I find nano particles of copper, and a clear solution. That's it. Now, if there is Cu 1+ remaining in the solution, I haven't checked.

Got experimental details, from some folks that claim to have the Cu1+ thing..... pretty well wired.

https://www.dimanregional.org/site/handlers/filedownload.ash...

Thanks for the link. I reproduced the lab experiment, at 1/12.5 mass and volume reduction. I got the results predicted by the link. I even used impure/iodized table salt (600mg) and 99+% root killer CuSO4.5H2O (800mg) with ascorbic acid (600mg). Instead of 50mL/solution I used 4mL. I shook the CuSO4 + NaCl solution, but did not shake the combined mixture of ascorbic acid in water + the salt solution in water. The combining of two solutions was done at room temperature 21.5C. I simply transferred ascorbic solution using an eye dropper into the copper solution. I saw it begin to turn colloidal immediately with the blue solution sinking to the bottom and turning green; so I photoed mid reaction; and then after transferring all the ascorbic acid, I capped the vial and left it for 10 minutes. There are three layers at the end (leftmost shot) in the vial; The expected white precipitate is on bottom. an unexpected dark green aqueous layer is on top of that; and then clear solution to the top of the vial.

The stoichiometry as shown in the paper is probably off a little bit, as the green layer does not precipitate even after a long time. But I replicated the ratio of chemicals in the lab paper, rather than work it out theoretically. ( EDIT: ideas were written after one hour ... see photo below for 16 hours later. Crystalization (not precipitation) can be seen on the glass of the vial in the region where the green layer still is. The color hasn't really changed, but it's much more transparent. The diffusion of liquids is amazingly slow in this reaction. The green band is mostly still visible at the bottom of the vial, with only a tiny amount of green color in the nearly clear fluid above it.)

You can see the color change from pure blue (leftmost picture SO4) to aqua-marine (2nd to left SO4 + Cl) because of the addition of table salt to the solution. So the color change mentioned in the lab worksheet is obvious (in spite of my blue-biased camera.); But both vials are still copper II ionization state.

I've repeated the same experiment using glucose, which is a reducing sugar with more hydrogen but the same oxygen's and carbons as ascorbic acid. ( C6H12O6 ). I used the same gram weight, which means I used slightly less moles of glucose (kitchen alchemy brand) than vitamin-C in the previous experiment (an error of 4g/mole, out of 180g/mol is ~2%); I still expected it to work with perhaps 2% more green solution left in the middle layer.

But an instant precipitation like happened with ascorbic acid, failed to occur. There are still three layers of fluid, but the color is still aqua-blue-green on bottom as if copper chloride/salt solution was not affected at all. A foggy colloidal layer exists in the middle; and clear sucrose solution on top. ( I'll post photos later, when the experiment is finished. )

I am letting it sit right now and seeing if it diffuses or precipitates before forcibly mixing it or attempting to heat it. I will probably begin warming it in the morning.

[Edited on 27-12-2017 by semiconductive]

[Edited on 28-12-2017 by semiconductive]

semiconductive - 27-12-2017 at 20:25

I've still been thinking about this carefully, and I still can't see how it can be true in a detailed sense.
All copper atoms act identically, as do copper ions. Therefore, if a copper ion separates from it's chloride in water (during "hydration" of copper+ chloride-) and thereafter doesn't interact ... then all copper atoms in the +2 oxidation state are colored purely by the copper ion's interaction with water because the other ion is nowhere near. Therefore, all copper +2 ions should have EXACTLY the same color dispersion in water because it's the water-copper interaction that determines the copper ion's hue.

Yet, the experiment I just did, and the Wikipedia article on Copper II prove the point that it has more than one color depending on which counter-ion is closest to copper ions. ( Sulphate or Chloride, or even concentration of Chloride. )

In my last experiment, the copper sulfate (A pure blue color) turns into a greenish-blue color upon addition of table salt. In both cases, it's still copper II (Cu++) ions that are present.

However, table salt by itself makes clear crystals and solutions -- therefore the ions of sodium and chlorine have no color at the concentrations used. If I just dilute blue copper sulfate with R.O. water ... it will only become paler blue (more transparent), but never green. The hue does not change (but the intensity/clarity does) when diluting cupric sulfate. ( Cu++SO4-- )

So, the negative charge on chloride ions has to be affecting the ground state of the electrons on copper in order to color it greenish; and for that to happen the chlorine ion's charges must be close enough to affect the electronic orbitals of the copper ions differently than does pure water.

Wikipedia also shows three vials of copper II chloride, and the color of the solution changes depending on the concentration of chloride.

https://en.wikipedia.org/wiki/Copper(II)_chloride

There's only two ways that I know this could happen based on Physics and atomic orbitals;
1) the chloride ions are physically close enough to the copper atoms to shift the electron orbital potentials.

2) The average electrical/back-"ground" potential of water changes depending on how many negative and positive ions are present; so that the color of ionization depends on how many chloride ions are present in the solution and not their location.

If the second one was the case, then if I stuck a wire into the ground from a battery ... and the other wire from a battery into a a copper chloride solution so as to change the "average" potential of the water; the hue ought to change. (It would make a neat liquid crystal replacement for a computer screen!!!) BUT ... It doesn't.

So, the conclusion I come to is that chloride ions are still localized near copper ions on average. I'm still guessing they must be within a very few water molecules diameter's distance of copper the majority of the time. eg: The chlorides must be close enough to change the average dipole moments (angles of water) significantly near the copper while overcoming intense thermal agitation of water molecules. Otherwise, it's impossible that copper II chloride's color wavelength and not just the intensity would change just because there are colorless chloride ions added to a solution of copper sulfate (in low concentration, it's not even 30 weight percent NaCl in my experiment!!! )

Table salt, Na+ Cl-, makes colorless and clear water-brine solution even when very strong.
(384 mg/mL , 1.2 g/mL density = 32% weight %).
However, HCl, (a different + ion) makes a pale yellow-green solution at that concentration.
The pale-ness could be an effect of how often the hydrogen interacts with chloride, so that darker color suggests ions which interact more. But I think the color wavelength change is a distinct measurement of which ions are interacting and how close they are to their counter during a "collision" or on average. ( eg: intensity = could be random collision effect alone, but wavelength change requires modulating an average distance between ion and counter-ion during the collisions/near misses themselves. )

So, I don't see how it's possible that the ion's are NOT interacting with their counter-ions extremely often (/on average) in order to affect their color. How do you come to the conclusion that they have "almost no interaction?" How do I objectively interpret "almost none" / use it in meaningful calculations to make verifiable predictions ?

[Edited on 29-12-2017 by semiconductive]

semiconductive - 29-12-2017 at 17:14

Even after heating to from 25C to 55C over two days, the solution is still aqua-marine.
The slight colloidal fog that formed immediately after transferring the glucose solution into the copper salt vial cleared up after an hour. Over the last two days, the copper solution has slowly diffused up into the glucose solution. However, even now only the top 5mm of solution looks very pale compared to the bottom. Glucose is obviously not going to precipitate CuCl like ascorbic acid, though I'm still not sure why. ( Edit: There actually is a very small amount of white precipitate on the bottom of the vial that showed up after shaking the vial.!!! )

I precision measured the starting substances and had an assistant double check with tare weighting the vials, measuring into a paper cone, and then adding chemicals and re-weighing. No mistakes were made. Mass in photos is +-1mg. It's a fairly precise repeat of the ascorbic acid experiment.

[Edited on 30-12-2017 by semiconductive]

semiconductive - 20-1-2018 at 22:04

The vials have been sitting now for weeks. The first picture is how each of them looked just before I removed all the water except a drop from on top. The second picture is after I removed the original aqueous solution from the yellow vial (ascorbic acid bath) and replaced it with water that was saturated under butane gas to drive off as much oxygen as possible. As you can see, the surface started to turn green a few days after the exchange in spite of the air being replaced by butane.

I plan on measuring the mass of the vials to get an amount of precipitate; but removing the water without oxidizing the precipitate is a problem. I'm waiting for some special teflon tubing and gas exchange membranes so I can try drying precipitates under butane gas. There is no sign that the white precipitate from the glucose solution is oxidizing to green. Its also not dissolving in water, so it's not glucose crystallized on the bottom. The present state of the vials is shown below; On the left two is the glucose precipitate and solution, the right two are ascorbic acid precipitate and remaining brown solution. I kept the solute before rinsing in the right vial of each pair. The color changes speak for themselves. The only source of oxygen getting to the T8.051 vial is from the water ionizing to H+ -OH. The top 0.25mm or so is now green. There is no green color change in the precipitate in the T8.267 vial.

[Edited on 21-1-2018 by semiconductive]

Rhodanide - 22-1-2018 at 10:50

From my experience, reacting (but not neutralizing) Basic Cu Carbonate with 31.45% HCl, aggressively boiling the solution with copper metal for 15-20 minutes, and pouring the resulting green-black liquid in to cold distilled water produces a lot of Cu (I) chloride precipitate. Personally, I hate the salt-reduction method. It's an unreliable method, and has never worked for me. The Copper-reduction method has worked beautifully for me five times now, and I plan to do a sixth run. I have an entire pill bottle filled to the brim with CuCl, wet with Ethanol. Only the top has oxidized, everything underneath is nice and pearly white!

[Edited on 22-1-2018 by Tetra]

DraconicAcid - 22-1-2018 at 11:48

Ions such as chloride or bromide will act as ligands to transition metals like copper; replacing water molecules as ligands will change the colour. Dilute solutions of copper(II) chloride will be as blue as solutions of non-coordinating ions like sulphate or nitrate. As you increase the concentration of chloride, you get greens and yellow. If you re-read the bit you quoted, you'll see that I was talking about compounds without transition metals. Sodium, potassium, cesium, calcium will never be coordinated by chloride or sulphate at reasonable concentrations.

DraconicAcid - 22-1-2018 at 11:55

Quote:

The molecular formula I learned for basic copper carbonate is: HO-Cu-CO3-Cu-OH
The carbonate CO3, has one double bonded oxygen, and two single bonded oxygens.

Cu-OH
|
O
|
C=O
|
O
|
Cu-OH

But as I pointed out, it's an ionic compound, not a molecule. The anions may at best be coordinated to the copper, but they aren't forming traditional covalent bonds.

Quote:

Three bonds puts oxygen in the -3 state .... !?

No. The number of bonds, covalent or coordinate, does not determine the oxidation state.

Quote:

It's cool looking, but I have no idea what kind of bonds it's trying to depict. Oxygen doesn't make three single bonds.

Ever hear of a hydronium ion?

Quote:

Hydrogen monoxide ?

Hydroxide ion.

color of copper solutions

wg48 - 22-1-2018 at 12:32

Here is a note that explains the color changes of copper ions.

Attachment: colcu bancroft1932.pdf (816kB)
This file has been downloaded 273 times

DraconicAcid - 22-1-2018 at 12:50

Here is a note that explains the color changes of copper ions.

What's the date on that? 1933? 1938? Their theoretical explanations and musing about the colour of anhydrous copper(II) ions is almost painful to read.

A copper(II) ion without any ligands will have no colour, because all the d orbitals will be degenerate. If you add ligands, they will stabilize some d orbitals and destabilize others, causing an energy difference between them. If an electron goes from a low-energy d orbital to a high energy one, it will absorb a photon of visible light, resulting in a colour.

The energy difference between the orbitals will depend on the kind of ligand (halide ion, water molecule, cyanide ion, amine) and the arrangement (tetrahedral, octahedral, square planar, distorted octahedron).

Octahedral arrangements (or distorted octahedral) of copper with oxygen-based ligands are almost invariably blue (usually pale). Nitrogen donors such as amines will usually form square planar or distorted octahedral with four amine ligands (sometimes two other ligands across from each other) which tend to be deep blue or purple.

Anionic ligands such as halides will generally give tetrahedral complexes even when water is also coordinated to the metal (such as [CuCl3(H2O)]-). These will often be green, or yellow (especially with chloride).

ETA: A better explanation = https://en.wikipedia.org/wiki/Crystal_field_theory

[Edited on 22-1-2018 by DraconicAcid]

semiconductive - 23-1-2018 at 19:33

Quote: Originally posted by Tetra

From my experience, reacting (but not neutralizing) Basic Cu Carbonate with 31.45% HCl, aggressively boiling the solution with copper metal for 15-20 minutes, and pouring the resulting green-black liquid in to cold distilled water produces a lot of Cu (I) chloride precipitate. Personally, I hate the salt-reduction method. It's an unreliable method, and has never worked for me. The Copper-reduction method has worked beautifully for me five times now, and I plan to do a sixth run. I have an entire pill bottle filled to the brim with CuCl, wet with Ethanol. Only the top has oxidized, everything underneath is nice and pearly white!

[Edited on 22-1-2018 by Tetra]

Thanks.

You're basically repeating Draconic Acid's suggestion on the first page. I intend to try that method also, as it's a worthy technique. However; The salt method did work for me, so I don't think it's unreliable so much as very sensitive to stoichiometry. I'm learning a lot by messing with the salts, but can understand why salt would fail for lots of people with limited equipment.

semiconductive - 23-1-2018 at 19:50

Ions such as chloride or bromide will act as ligands to transition metals like copper; replacing water molecules as ligands will change the colour. ... If you re-read the bit you quoted, you'll see that I was talking about compounds without transition metals. Sodium, potassium, cesium, calcium will never be coordinated by chloride or sulphate at reasonable concentrations.

I'm a newbie; I would never have made the connection between the type of ions and replacing of water molecules based on the type of ion on my own. My background is, again, semiconductor physics in the solid (crystalline) state. I only have 100 and 200 level chemestry courses in my background, and no o-chem. Your clarification is very much appreciated. I'm beginning to understand a few other things you said, now.

Transition metals are special because of D orbitals, so there are up to 10 electrons to fill the D shell. Copper is the 9th electron in the period, so it is only one electron shy of a full shell. Zinc, the next metal, would be questionable as a transition metal because it's shell is full an unreactive.

Still, the problem with the qualitative ideas you are giving me is that I have no experience to decide when something is "normal" or "high" concentration. Just as I can't know how far apart ions will stay (on average) in a liquid -- regardless of whether they are transition metals or not.

Take for example, hydrogen. It's in the first two columns of the periodic table ... so that means 1S and 2S orbitals are on the outside. According to what you have said, it's not something that would be coordinated by sulfate or Chloride; for it's in the same S orbital columns of the periodic table as sodium and calcium. However, sodium chloride is clear at very high concntrations (even saturation); but Hydrogen Chloride is greenish in test tube solution. So, Sodium vs. Hydrogen changes the color of the solution. But, hydrogen is part of water ... and there are [H+] [OH-] ions floating around in water ... so again, if Na was not close to Cl ... then the hydrogen in water should make NaCl solution green ... but it's clear.

There are often exceptions to the things you have been stating in an almost absolute sense. It's hard for me to track without quantitative ways to calculate outcomes why the things you say are true, and in what sense. You have a lot of experience and easily recall exceptions, but I don't. That weakness of mine is obviously a big part of my troubles. I don't have a way to get a "feel" for what is normal and concentrated... what is "close" together in a solution, and what is far apart ... etc. To me, they are all atoms with often invisible and complicated interactions going on in solution. It's extremely difficult for me to reason to a conclusion in such a situation. I do appreciate your patience.

DraconicAcid - 23-1-2018 at 20:41

Quote:

Yes, zinc is often not considered to be a transition metal, and because it always has full d orbitals, its compounds are never coloured unless a) the anion is coloured (such as zinc permanganate) or b) the ion is very easily reduced (which gives colour due to charge transfer- this would be the case for lead(II) iodide, but I can't think of an example for zinc).

Quote:

Still, the problem with the qualitative ideas you are giving me is that I have no experience to decide when something is "normal" or "high" concentration.

That depends on the ions in question. Cyanide is an excellent ligand, and will coordinate even at low concentrations. Chloride may coordinate to copper at *waves hands* say, 1 mol/L, but it won't coordinate to nickel or manganese unless at a much higher concentration. Choride won't coordinate appreciably to iron(II), but will coordinate readily to iron(III). Nitrate, perchlorate and sulphate won't coordinate unless the solution is so concentrated that it no longer really qualifies as an aqueous solution.

Quote:

Just as I can't know how far apart ions will stay (on average) in a liquid -- regardless of whether they are transition metals or not.

On average, the tendency for a metal ion to be coordinated by, or at least pair up with an anion is determined by the radius of its charge to its radius. A small and/or highly charged ion will be coordinated much more readily than a large one with a low charge. This is why the alkali metals don't do coordination chemistry.

Quote:

Take for example, hydrogen. It's in the first two columns of the periodic table ... so that means 1S and 2S orbitals are on the outside. According to what you have said, it's not something that would be coordinated by sulfate or Chloride; for it's in the same S orbital columns of the periodic table as sodium and calcium. However, sodium chloride is clear at very high concntrations (even saturation); but Hydrogen Chloride is greenish in test tube solution. So, Sodium vs. Hydrogen changes the color of the solution. But, hydrogen is part of water ... and there are [H+] [OH-] ions floating around in water ... so again, if Na was not close to Cl ... then the hydrogen in water should make NaCl solution green ... but it's clear.

Hydrogen isn't an alkali metal, though- it's a non-metal, despite being in the same column as sodium. The H+ cation doesn't have *any* filled orbitals, so it would be orders of magnitude smaller than any ion. Thus, it *always* forms a covalent bond with something- an anion, or a water molecule (to make the hydronium ion H3O+). That's why hydrogen chloride is a molecular gas at room temperature, and sodium chloride is an ionic solid. If HCl is greenish, it would be due to electrons moving around in molecular orbitals, not the d orbitals responsible for transition metal compounds.

Quote:

There are often exceptions to the things you have been stating in an almost absolute sense.

I'm trying to keep things simple.

Quote:

I do appreciate your patience.

I'm glad I still seem patient- my students think I've been snarling at them lately. But I hope my comments have been helpful.

semiconductive - 24-1-2018 at 13:44

Quote: Originally posted by Tetra

In a perpetual search for nitric acid.

I came across a youtube video, where a person put a jacob's ladder with copper electrodes inside a fish bowl with a sealed lid. The strength of the acid was surprising after an overnight.
All it needs is an old broken microwave's transformer.

https://ru-clip.com/video/ep23ds4cZs4/nitric-acid-from-thin-...

semiconductive - 24-1-2018 at 14:09

Here is a note that explains the color changes of copper ions.

Thanks. That's very helpful in an Amateur's ability level. There were several anecdotes that were useful in me forming a stronger opinion about what's really going on, intuitively.

I had to laugh though, page 1064: "7. Copper Oxide is blue and not black"
That's quite possible as dark blue and black are hard to distinguish by eye.
I know from expeience that copper I oxide looks reddish, and copper II oxide looks black to the naked eye.

But even the author, after realizing that the (II) oxide ought to be blue, said a few pages later: 1068: This is apparently due to to formation of undissolved, black, cupric oxide.

So there are some points which are ambiguous because of human sight.

I've purchased a 1000 line/mm diffraction grating and have been planning to build a spectrometer capable of visible and infrared measurements down to 10 micron. Unfortunately, plastic of diffraction gratings are not good to infrared wavelengths used in transmission mode; and to get it to work in reflection mode, I need to silver the back side of the film. I tried talking with the local high school chem teacher who has his students build a crude spectrometer for visible wavelengths using an LED, and although he's excited about the idea of a wide range spectrometer for low cost, he doesn't have the time to help me (wife needs attention kind of problems...!). So, it won't be until one of my sons is in his Chemistry class next year that I have an excuse to pursue an advanced spectrometer with him. But that's what is really needed to fine tune the experients being conducted on color. Thermal agitation ought to change the colors of the ions slightly, not enough to be seen by the human eye or a simple camera, but something that could be measured by a simple spectrometer. That would be an invaluble tool in figuring out what is really going on in aqueous chemistry.

Again, I appreciate the link. I learned a lot.

It's fairly obvious that the solute changes affected the color based only on molecules that were in intimate contact or proximity to the copper. Not more than one molecules distance away. That means color is affected more by the number of water molecules, vs. counter IONS displacing water molecules than even electric fields from nearby ions. eg: the color of copper is affected more by molecules DISplacing water, than by their interaction with the copper itself.

The author's distinction between opacity (clarity), and color, is not always clearly stated; but that's one of the major reasoning tools to figuring out what is really going on. The shade of color vs. intensity, allows one to crudely distingusih between degree of ionization, and type/cause of ionization/coordination.

Cuprous ions obviously have the energy potential for producing the red-yellow range, while Cupric ions have energy levels compatible with the green-blue range of colors.

Browns and blacks, being combinations of colors, need to be identified with a spectrometer and not the human eye for an accurate assesment of the combinations of greens, blues and reds which are treuly present..

[Edited on 24-1-2018 by semiconductive]

DraconicAcid - 24-1-2018 at 14:16

The article is good at being descriptive, but suffers from being written prior to the development of crystal field theory. The colour of the copper(II) ion is not determined by the number of water molecules present, but the number and kind of ligands attached to the copper, of which water is only one kind.

Σldritch - 24-1-2018 at 14:23

Reducing sugars are reducing because they can become aldehydes in solution. Ascorbic acid is oxidised to form a radical. Why one seems to reduce Copper (II) to Copper (I) Chloride in these conditions i do not know. At least it is a clue.

semiconductive - 24-1-2018 at 17:01

Quote:

[/rquote]

Beautiful; and you're proving my point about the deep nature of your experience and empirical knowledge which a newbie lacks.

Quote:

Just as I can't know how far apart ions will stay (on average) in a liquid -- regardless of whether they are transition metals or not.

I see, and that is the most useful thing you have told me so far in this thread.

Thank you!

Quote:

Hydrogen isn't an alkali metal, though- it's a non-metal, despite being in the same column as sodium.

Yes, but think about that for a moment. Metals are classified as metals from their physical properties when crystallized; eg: conductivity, and related reflectivity (shiny!). For it's the conductivity, when maxwells equations are solved that has a lot to do with succesfully making a mirror. There are different ways to get conductivity, but the idea of metal .. is one where in the crystal state the molecular orbitals overlap and are only partially filled. That means the topmost electrons are free to travel because they can "jump" to nearby energy states that are vacant. In contrast, non conductors have a filled valence band ... but there are no empty states nearby for electrons to "jump" to. It takes significant energy, on the order of a photon of visible light, to promote electrons in a non conductor.

No metal is really acting like a "metal" when ionized in water solution. The bands surrounding the ions are at most molecularly coupled to a handful of very nearby water molecules. So, the important distinction has notihing really to do with being a metal and a "conduction" band ... The important distinction is how a particular atom molecularly hybridizes it's orbitals with water or other solvents, and why.

Quote:

The H+ cation doesn't have *any* filled orbitals, so it would be orders of magnitude smaller than any ion. ... If HCl is greenish, it would be due to electrons moving around in molecular orbitals, not the d orbitals responsible for transition metal compounds.

Of course. I agree, that's why I pointed out 1S and 2S orbitals as being the outer unfilled orbitals involved in bonding. It's just that those are the same shaped orbitals as the other alkalai metals, and ought to behave similarly. In the case of hydrogen, then, as an exception it DOES matter that the chloride atom is floating nearby or at very least displacing water from contact with the hydrogen to change the color. The exact color will depend on the quantum states which are stable in whatever solvent the hydrogen is in and the counter-ion which is near it (if any).

Quote:

I'm glad I still seem patient- my students think I've been snarling at them lately. But I hope my comments have been helpful.

You might try starting your sentences with a different word than "no." Whether you are correcting a small mistake or a big one, the word "no" can make a student feel like a dog or baby in a family. Once the adrenaline rushes, their ability to reason is often diminished. I realize there are some bone headed students that you have to get their attention before they will even listed ... ADHD people ... but for your brighter students, you might try using a less absolute opener to your corrections; eg: "not quite", or "have you thought of..."

Your comments are a very steep learning curve. In the past, I would say they have not been extremely helpful ... but that's not to say they won't prove more useful in the future. For me, I still haven't seen copper oxide dissolve in ammonia water. I have ONLY seen it dissolve in ammonia water exposed to air and expecially carbon dioxide. So, it's not that you are necessarily wrong ... but you've given me a contradiction to my experience that I need to test and quantify. If something dissolves only a handful of atoms out of a gram ... that might be what you mean; but I have no way to grasp these things without testing them.

I'm also fighting several sources of error. My scale, for example, has a hysterisses effect. I've been testing boyancy issues, water absorbtion of air, and other reasons for my scale giving nonstandard results. BUt, after the past few days I realized that the mechanism inside is a strain guage on a metallic bar; that guage deflects more or less depending on it's immediate past history and possibly even magnetic state. The software is pretty good at hiding it ... but it's still giving me errors of up to 4 milligrams. So, the experiments I did in this thread need to be redone once I've figured out ways to degauss the strain guage during operation and weighing of mass.

It's all the little things that keep biting me in the ass, that are making it difficult to isolate individual issues as is required in the scientific method to "test one thing at a time" while keeping all others constant. It's not your comments, so much as the time it takes to isolate the issues.

[Edited on 25-1-2018 by semiconductive]

semiconductive - 24-1-2018 at 17:30

ETA: A better explanation = https://en.wikipedia.org/wiki/Crystal_field_theory

[Edited on 22-1-2018 by DraconicAcid]

DraconicAcid, how's your quantum mechanics?

That article gets into the nitty gritty and I've been studying these issues of magnetic spin coupling and non-degeneracy for semiconductors for over six months now.

For example, here's a thread I opened on electronic's stack exchange:
https://electronics.stackexchange.com/questions/309097/semic...

The graph tells a lot; for I'm comparing experiment to theory and you can see the graphs don't quite match. I've spent six months buying master's thesis, asking questions on physics forums, etc. and basically; nobody knows how to solve these issues. The issue is spin-orbit coupling, and landau level effects.

Bascially, the author's of the original article I tried to reproduce used a wrong equation to compute the effective mass of electrons. ( Which even in molecular orbitals can be used to test them for the ability to make certain colors. )

Your crystal theory link makes an interesting assumption that is directly related to band theory (molecular orbitals when statistically many atoms are nearby) and the magnetic fields. Unfortunately it will be wrong in many cases; but if the crystal theory is right, then semiconductor band theory MUST be wrong in a large number of cases.

I have been asking people about the details of QM in magnetic fields... but the online physics forums are basically worthless. I end up solving everything myself with no guidance ... and it's taken months to realize and prove where the problem lies. I have experimental evidence to back up my conclusions ... but seeing the crystal theory paper makes me groan because there's another subtle issue in it.

https://www.physicsforums.com/threads/cyclotron-resonance-an...

Do you have the background and time to dig into the reasons why the Aufbau principle is clearly being violated in the "high spin" drawing of the link you gave? The author only mentions that the "Hunds rule" is not followed (There are more than one rule called, Hunds...)

There is clearly a spin, or angular momentum, and orbital interaction issue similar to zeeman splitting and landau levels in band theory at the heart of crystal theory.

State of the art semiconductor physics as taught in colleges at the masters level, is wrong in many cases because they ignore how much the magnetic fields of cyclotron measurements affect the energy of electrons. That's why I linked stack exchance, was to show how far off the standard master's levels analysis of semiconductors is.

The reason is that the analysis is usually semi-classical.
In classical physics, magnetic fields can only change the direction but not energy of an electron. However, in QUANTUM physics the Heisenburg uncertainty basically requires that even a magnetic field MUST affect the energy of an electron in an orbit.

Eariler, I ingored your offer to work out mass action law in terms of thermodynamics; becasue the first page of my thermodynamics book says "The laws of thermodynamics are empirical laws..." ... so there's no point in doing that to prove mass action which my teacher said was "an empirical law". But if you are able to work QM problems with me, that would be very helpful in understanding crystal theory !

[Edited on 25-1-2018 by semiconductive]

DraconicAcid - 24-1-2018 at 17:37

Quote:

Yes, but think about that for a moment. Metals are classified as metals from their physical properties when crystallized; eg: conductivity, and related reflectivity (shiny!).

No metal is really acting like a "metal" when ionized in water solution. The bands surrounding the ions are at most molecularly coupled to a handful of very nearby water molecules. So, the important distinction has notihing really to do with being a metal and a "conduction" band ... The important distinction is how a particular atom molecularly hybridizes it's orbitals with water or other solvents, and why.

That may be the physicist's definition of a metal, but a chemist defines a metal as an element that tends to lose electrons to form cations when it reacts with a nonmetal. Hydrogen, unlike sodium, lithium, etc, does not form simple cations, and always forms either molecular compounds, or becomes part of a covalently bonded polyatomic cation. So sodium, lithium, potassium etc. all act like typical metal ions in solution, whereas hydrogen shows very different behavior.

Quote:

In the case of hydrogen, it uses its 1s orbital for bonding. For the alkali metals, apart from very rare compounds such as methyllithium, their bonding is almost purely ionic, and they don't use orbitals for that- just cation attracting anion.

[Edited on 25-1-2018 by DraconicAcid]

semiconductive - 24-1-2018 at 18:10

My copper sugar experiments are only in water; so that's the part of the article I was focusing on; eg: that which applies to the case at hand.

To be fair to the author, he did not make a mistake that was reducing the issue to water alone. Neither did I intend to imply that only water was involved in color making. The paper clearly states that the hue of "blue" changes when ammonia is the solvent instead of water. Therefore, even this ancient author gives an example proving they knew that color depends on kind of ligands and well as number.

semiconductive - 24-1-2018 at 18:23

Quote:

OK. Another background issue. Tomato, Tamatoe.

Quote:

The difference between an ionic bond and a covalent bond, quantum mechainically, is usually one of degree and not kind. Every time atoms come together there are filled orbitals/ stationary states, and unfilled orbitals/stationary states. AKA, the so called "bonding" and "anti-bonding" orbitals. But any actual motion of a wave-particle requires the superposition of two or more states. Therefore, any time atoms are in thermal motion .. multiple quantum states are involved.

Some people think that when an atom vibrates due to thermal agitation, the electrons are promoted for brief periods of time into the lowest anti-bonding orbital. Others think the orbitals split. So, it's an over-simplification to say that all the alkali-metals are purely in an ionic state. Atoms somehow share electrons in proportion to the vibrations and tunneling occurs a fair number of times a second. I mean, E=hf can be applied to the bonding energy of two atoms as well as to the wavelength of a moving electron. So, if you tell me a bonding energy -- I should be able to compute a vibration frequency for the two atoms around a center of mass, or "barycenter." as relativity buffs like to correct me when I use the older term; and show how often the electrons go around the counter-ion. It's a non-zero probablity.

Is this why the drawing you showed me from wikipedia has hydroxide (OH) ions with no bonds being shown to any other nearby atom? They are making some kind of hard distinction between ionic bond and covalent bond???

I've heard and seen chemistry where resonance effects are going on. In older texts, that was usually shown by a dotted line where a bond switched from one atom to another. Indeeed, in the pictures of Citric acid one such bond is shown between the two OH radicals. The same type of bond is NOT shown in the case of glucose.

But I have never seen someone draw an oxygen atom having three bonds ... like the picture you showed me from Wikipedia. And no, I've never heard of the ion you asked me about.

We just treated ph as H+ and OH- in solution. Since Oxygen has already been satisfied with two hydrogens donating electrons to the shells; I would think a third hydrogen was merely a resonance effect where it time shares bonding to the oxygen with other hydrogens. I don't see why you couldn't have a fourth hydrogen, making a H4O++ ion, though it would be even less stable than the so called hydronium ion.

But I don't see what concept is important when the hydronium ion idea was made. In fact, every time a water molecule with an extra hydrogen tagging along randomly bumps into another water molecule -- it's equally likely that a hydrogen will transfer from one water molecule to the next. The "thrid" hydrogen isn't stuck to any given hydronium molecule electrostatically more than another water molecule. I would think it's bond energy would be indentical with all water molecules ... shrug ...so it effectively wanders freely (or makes the whole hydronium "molecule" unstable so that each of the hydrogens is temporally equally likely to wander away from the hydronium ion). I don't see how it's more advantageous than earlier resonance line drawings of water...

[Edited on 25-1-2018 by semiconductive]

[Edited on 25-1-2018 by semiconductive]

DraconicAcid - 24-1-2018 at 18:23

To be fair to the author, he did not make a mistake that was reducing the issue to water alone. Neither did I intend to imply that only water was involved in color making. The paper clearly states that the hue of "blue" changes when ammonia is the solvent instead of water. Therefore, even this ancient author gives an example proving they knew that color depends on kind of ligands and well as number.

When he talks about ammonia, he does say this, but when he's talking about various chlorides (i.e., copper(II) chloride and mixed salts with cadmium and other chlorides), he does dismiss the chlorides and speaks only of the number of water molecules. I wasn't suggesting that you were implying anything.

semiconductive - 24-1-2018 at 18:35

Quote: Originally posted by Σldritch

Reducing sugars are reducing because they can become aldehydes in solution. Ascorbic acid is oxidised to form a radical. Why one seems to reduce Copper (II) to Copper (I) Chloride in these conditions i do not know. At least it is a clue.

Thanks Eldritch. Yeah, it's a real stumper. Do you have any resources/online links where the aldehyde connection is explained? The wikipedia stoichiometry for the reduction of copper shows only that the net effect is the two hydrogens are lost (presumably to form water). How an aldehyde in citric acid would cause the two hydrogens to be lost is a mystery to me. I would appreciate any links to clues.

[Edited on 25-1-2018 by semiconductive]

[Edited on 25-1-2018 by semiconductive]

DraconicAcid - 24-1-2018 at 18:54

Is this why the drawing you showed me from wikipedia has hydroxide (OH) ions with no bonds being shown to any other nearby atom? They are making some kind of hard distinction between ionic bond and covalent bond???

I've heard and seen chemistry where resonance effects are going on. In older texts, that was usually shown by a dotted line where a bond switched from one atom to another. Indeeed, in the pictures of Citric acid one such bond is shown between the two OH radicals. The same type of bond is NOT shown in the case of glucose.

But I have never seen someone draw an oxygen atom having three bonds ... like the picture you showed me from Wikipedia. And no, I've never heard of the ion you asked me about.

I skip the quantum mechanical parts because, for my own sake, I'm sticking to the simpler models.

The hydroxide shown in the picture isn't shown as bonded to anything, I think, only because it's bonded to atoms that are outside of that particular unit cell. They aren't making a hard distinction between ionic and covalent bonds- they show a bond whenever the two atoms are within a certain distance.

A hydronium ion is the H3O+ ion that you get in aqueous solutions of acid. Any time you have a water molecule acting as a ligand to a metal, you will also have an oxygen with three bonds.

Check out https://www.ccdc.cam.ac.uk/structures/?

A lot of the structures are too complex to be of interest to me, but if you search for "diaqua" and "copper" you can find some nice ones (and view them as the unit cell, or a collection of unit cells. They don't have basic copper carbonate, though).

semiconductive - 24-1-2018 at 20:17

in the paper, he noted that anhydrous cupric chloride was "yellow-brown" (P1066); so naturally, he's going to talk about the change in color due to the number of molecules added for hydration. I don't see how you come to the conclusion that he "dismisses" chlorine? The chlorine obviously doesn't cause the bluish color, on it's own; it's the water molecules that shift the color to blue.

I didn't pay attention to the mixed salts, because I'm not working with them yet. It's information overload. But, if we know water is sufficient to cause an ion to be a particular color ... and changing anions does not change the color, SO4, Cl, then why would he pay attention to the particular anion? Wouldn't he be justified in assuming it plays a minor role and maybe shifts the color an imperceptable amount? With the anions in scope, the water is the one who's bonding interaction (with some SO4,Cl, or other anion assumed) and the copper that causes the distinct effect?

Which statement, on what page, are you singling out? I'm missing something.

semiconductive - 24-1-2018 at 20:37

They aren't making a hard distinction between ionic and covalent bonds- they show a bond whenever the two atoms are within a certain distance.

I see. I've always understood bond sticks to represent the sharing of electrons to fill the orbitals; eg: octet rule and so forth. So, it's very foreign to me to see a stick where there should be a dotted line for resonance. It's been 25 years, but the only other representation I recall seeing is one where dots were placed between atoms to represent the number of electrons shared ... lewis diagram, maybe? Sometimes we would x out one of the dots to show it didn't come from the adjacent atom ... but it's been too long.

The wikipedia drawing was just very confusing to me. The diagram I drew was merely to show an accounting of one configuration whiere all the orbitals were properly filled by shared electrons. It's not impossible for there to be resonances not shown in my diagrams; but at least the diagrams I showed satisfy the electron sharing requirements. I don't see why they are intrinsically impossible.

DraconicAcid - 24-1-2018 at 23:17

The wikipedia drawing was just very confusing to me. The diagram I drew was merely to show an accounting of one configuration whiere all the orbitals were properly filled by shared electrons. It's not impossible for there to be resonances not shown in my diagrams; but at least the diagrams I showed satisfy the electron sharing requirements. I don't see why they are intrinsically impossible.

Because you've drawn it as if it were a simple molecule with copper having two bonds. The bonds with copper are coordinate bonds rather than classical covalent bonds. Cu isn't going to need two bonds to satisfy its valence- it's going to have a coordination number of four or six in most compounds.

Look- when you have sodium chloride, you don't have distinct Na-Cl molecules, and if you draw it as Na-Cl, you'll be told that it's wrong. In solid sodium chloride, you have a sodium ion surrounded by, and equally attracted to, six chloride anions. It's not bonded to any one chloride more than it is the other five close to it.

Here's another image of the structure of basic copper(II) carbonate:

Each carbonate bridges four copper atoms; each hydroxide bridges two of them (ignore how the structure shows each carbonate having two single bonds and one double bond to carbon- that's just the software requiring a full octet, even though the three C-O bonds are equivalent). Each copper is square planar, so there are sheets stacked on top of each other.

wg48 - 25-1-2018 at 02:29

I've purchased a 1000 line/mm diffraction grating and have been planning to build a spectrometer capable of visible and infrared measurements down to 10 micron. Unfortunately, plastic of diffraction gratings are not good to infrared wavelengths used in transmission mode; and to get it to work in reflection mode, I need to silver the back side of the film. I tried talking with the local high school chem teacher who has his students build a crude spectrometer for visible wavelengths using an LED, and although he's excited about the idea of a wide range spectrometer for low cost, he doesn't have the time to help me (wife needs attention kind of problems...!). So, it won't be until one of my sons is in his Chemistry class next year that I have an excuse to pursue an advanced spectrometer with him. But that's what is really needed to fine tune the experients being conducted on color. Thermal agitation ought to change the colors of the ions slightly, not enough to be seen by the human eye or a simple camera, but something that could be measured by a simple spectrometer. That would be an invaluble tool in figuring out what is really going on in aqueous chemistry.

Again, I appreciate the link. I learned a lot.

[Edited on 24-1-2018 by semiconductive]

If the problem of using the plastic grating at IR wave lengths in transmission mode is the absorption by the plastic, then silvering the rear will mean the IR must pass through it twice so the absorption would be greater. Perhaps you could silver the front then copper or nickel plate it sufficiently thick that it can be peeled from the grating.

I liked the note I posted because it discussed the colors of the various copper compounds while giving the chain of reasoning as to how the conclusions was reached. A more practical approuch than many more theory orientated books.

Your were welcome, I am glad it helped you.

semiconductive - 25-1-2018 at 13:20