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Author: Subject: Reagent Production by Electrolysis - a few ideas.
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[*] posted on 14-10-2009 at 16:04
Reagent Production by Electrolysis - a few ideas.

Electrolysis often provides a method of producing reagents more easily than standard chemical methods.These processes are always slower and often easier to control. Here are a few examples.If I haven’t tried any of the following I attach an asterisk (*); D indicates a divided cell, U undivided (without diaphragm).

First, concerning ELECTRODES:
We are interested either in metals or conductive substances that are (a) either inert under electrolysis, or (b) those that are consumed – the latter if we want to produce their salts.

Few cathodes react with the hydrogen nearly always produced. Even if they produce hydrogen from dilute acids they can still be used - provided they can be plated. Thus although iron produces hydrogen with dilute acid, it can frequently be used in mildly acidic conditions provided current is flowing in the cell. I use it with NaHSO4 with pH ~ 1.5. Any Fe dissolved either replates in acidic or forms insoluble hydroxide in alkaline solution. Typical cathodes are Fe, Ni, stainless steel (SS), C, Pb, Cu, Sn, Hg. Other metals (e,g, Ti (*)) are also used. Best is Pt but now unaffordable at more than $1200/oz.

Several of these occlude hydrogen, esp. Ni. I have noticed that Cu will form a dark scale as a cathode, which is CuH (AFAIK).

Nearly all these metals will oxidize as anodes in suitable acidic solutions to give metal ions in solution. Exceptions are Pt or when a higher oxide stable in acid is formed, e.g. PbO2. Deposited PbO2 anodes are the subject of a long thread on this forum (q.v) and elsewhere . Pb and PbO2 are very useful. If oxygen production is avoided, carbon, preferably graphite, can be used. The halogens attack most metals including Ag, Au, even Pt to some extent – so use carbon preferably.

In alkaline solution the iron metals Fe, Co, Ni form a thin coat of conductive higher oxides (believed to be Fe3O4, Co2O3, Ni2O3, - maybe NiO2, or CoO2) and can be used to evolve oxygen as efficiently as Pt does in acidic solution. Alkaline anodic processes should use these if the metal is not to be consumed. SS is also sometimes useful.

PRODUCING ACIDS: I previously posted a typical process in detail, for obtaining H2SO4 from sodium bisulphate. MgSO4 and FeSO4 are also satisfactory for this production (D) of H2SO4. The first produces a cloud of Mg(OH)2 at the cathode (Fe); the second deposits iron.

NITRIC acid is difficult to produce electrolytically from nitrates (except on a Pt anode(*)). All common metal anodes are oxidized and dissolved. Carbon can be tried but also oxidizes - the acid concentration must be kept quite low. Unlike sulphuric acid HNO3 is a strong oxidant even when comparatively dilute. It might be worth a try with a massive deposited PbO2 anode (*). Lead itself is converted to nitrate.

HCl – (D) the anode product is usually a chloride of the anode metal or, with carbon or PbO2 anode in concentrated electrolyte, Cl2 gas, and a mix of HCl and HClO. Similarly for other halogens. Note that both I2 and Br2 dissolve in KI and KBr solution to produce complex ions KX3-. Cl2 doesn’t.

(*)Lead or PbO2 covered lead, (D), should work with phosphates to give H3PO4 as this acid is non-oxidizing.

(*)Organics tend to oxidize at the anode but some monobasic aliphatic and aromatics might survive. However, salts of these are probably less common than the acids themselves. Oxalic acid would probably be oxidized to CO2.

(D)By using a conc. alkali metal salt solution in the anode compartment, Pb anode and Fe, Ni, or C cathode, (with initially a very dilute solution of the same salt in the cathode cell), the process produces a solution of alkali hydroxide at the cathode. The anode cell becomes acidic. For instance, if Na2SO4 is used, at some point the anode will contain effectively NaHSO4. Depleting beyond this point gives a low yield because H+ ions are far more mobile than Na+ ions. Hence stop once pH is less than 1.5 so that Na+ ions predominate.
The anode contents can then be used later in the cathode compartment for efficient H2SO4 production – until they become alkaline.. This way tolerably pure dilute acid and hydroxide can be gotten from a sulphate. Neither nitrate nor chloride can be used this way because of anode erosion and subsequent heavy metal intrusion.

Example: CuSO4 from Cu and NaHSO4:
Suppose CuSO4 is desired from scrap copper. The standard lab method oxidizes the metal with conc. H2SO4 using considerable heat. Noxious fumes of SO2 are produced and a mess of sulphate, unreacted H2SO4 and Cu, and often black CuO or CuS are left. Several recrystallizations are needed to get rid of excess acid.

The electrolytic method (D) uses a Cu anode, with slightly acidulated water to render it conductive, and a cathode (Fe, C, Ni etc) compartment of saturated sulphate solution, preferably NaHSO4 or dilute H2SO4 (i.e, acidic). The anode cell becomes acidic and SO4++ ions oxidize the copper and react with it to produce CuSO4 (properly, Cu++ ions) without any gas production.
In a similar way PbNO3 can be made from Pb using NaNO3 in the cathode cell, without noxious fumes of NO, NO2 or the requirement for HNO3.

Any metal M that forms a soluble salt MX with an acid HX can be so treated, e.g. M=Ag, Mg, Al, Pb, Mn, Fe, Ni, Co and X= SO4, Cl, NO3, PO4 etc; the useful cases are those where conc. acid has little effect (eg, HCl on Cu), the acid is unavailable or the process produces noxious fumes. Note, however, if the metal produces a higher oxide insoluble in HX (like Pb in H2SO4) that oxide and HX may be produced instead.

You can produce chlorates this way from KClO3 – but remember that many chlorates are dangerously unstable if dried.

With a Cu anode, (D) using saturated NaCl solution as catholyte and NaCl or HCl in the anolyte, instead of the expected CuCl2 passing into anode solution, white CuCl is also precipitated. I suspect the same happens with KBr or KI, giving CuBr or CuI.(*)

Note 1: once metal ions are present in the anolyte, they will pass through the diaphragm to the cathode cell and compete with H+ ions. The high mobility of H+ ions makes the process reasonably efficient if the anode cell is kept somewhat acidic.

Note 2: nitrates are reduced simultaneously at the cathode to nitrites and ammonia, and possibly also hydroxylamine (NH2OH) so the nitrate case is less efficient.

HEAVY METAL HYDROXIDES: (U) These are mainly insoluble except for thallium(I). If the same setup is as the last is used without a diaphragm, the soluble salt produced at the anode will react with OH- ions produced at the cathode and be precipitated as hydroxide, using an neutral alkaline metal salt as electrolyte eg Na2SO4 (but not NaHSO4). The hydroxide can then be dissolved in a dilute acid to give a desired salt. The cathode electrolyte is not consumed unless an insoluble salt is produced with it having a solubility product less than the desired hydroxide. The OH- ions come from water.

OXIDATION OR REDUCTION: Anodic oxidation and cathodic reduction, chemically speaking, is always the action of electrolysis in aqueous solution: Oxidation being the loss of electrons, reduction the gain. Often poorly- or un-ionized substances added to the electrolyte can be reduced or oxidized. Thus naphthalene suspended in H2SO4 in the anode compartment (D), with H2SO4 in the cathode, is said to be oxidized to phthalic acid(*) using a PbO2 anode. I have tried isopropyl alcohol to acetone & ethyl alcohol to acetic acid this way. Another example is said to be CH3OH to formaldehyde (*)..

Anodic oxidation of Cl- ions to ClO- and thence to ClO3- is well known (U, with C or (better) PbO2 anode) – see copious posts on this forum and elsewhere. ClO3- to ClO4- is more difficult (U, with PbO2 or Pt anode). In both cases chromate or NaF is often added to prevent simultaneous cathodic reduction. Why NaF works is a bit of a mystery – perhaps transient F atoms are produced that promote the anodic oxidation.

A divided cell is sometimes helpful to avoid cathodic reduction but in the chlorate cell the cathode product must be available at the anode and hence the cell must be undivided unless the catholyte is circulated to the anode.

Many metal ions can be reduced to metal at the cathode. The conditions for this demand either (a) the metal/ion Standard Electromotive potential (SEP) is negative – i.e. the metal is electronegative to H2. (e.g. Cu); or (b) the overvoltage of hydrogen on the metal exceeds the SEP if positive. Metals as reactive as Zn or Mn (SEP ~ +1v) can be deposited. Even Na and K can be separated if a mercury cathode is used with a strong chloride electrolyte, because the SEP of Hg/Na amalgam is of the order of 1v instead of 2.7v for Na metal. If NH4Cl is used as electrolyte with a Hg cathode, the mysterious ‘ammonium amalgam’ can be made, but watch out for the dangerously explosive NCl3 which can be produced under certain conditions (high concentration).

For production of metallic sodium from a fused salt see the posting by Len1 on the forum.

Metal deposits vary from powder, crystalline (E.g. Cd or Ag) to somewhat coherent plating. (Proper metal plating is a distinct art). Usually, if simple salts are used, e.g. sulphates, the coat can be stripped off, say, a C cathode. Generally strong, acidic salt solutions work best. Current density should be such than little or no H2 gas is given off.

Metals that have a tendency to produce a higher acidic oxide (Mn, Cr, V, etc) can often be anodically oxidized to permanganate, etc. by electrolysis in alkaline electrolyte. Solid MnO2 (which is conductive) , or even powder contained in a glass cloth bag around a C, Fe or Ni anode, can be oxidized to permanganate in alkaline solution (KOH). It is claimed that a slurry of MnO2 can also be (* see Permanganate thread, esp. Xenoid’s posts). K2Cr2O7 can be produced from Cr/Fe alloys or Cr+++ ions by anodic oxidation in KOH. Perhaps a high chrome stainless steel might work?(*). Manganate made from fusing KOH, KNO3, and MnO2 can be oxidized to permanganate. (I can get KOH to work but am never very successful with NaOH for some reason – chlorate works better than nitrate).

(*)Many organic substances can be reduced at the cathode – do a literature search if interested.

Persulphates are said to require Pt cathode and anode, (U), sulphate salt or acid, and temperatures around 5C or less, with high current density ( ~0.5 a/cm^2)(*). Chromate or HF is also suggested to avoid reduction at cathode. K2S2O8 preferred because of low solubility. Persulphates are rather unstable so crystallization in the cell is best. At extremely high current density O3 is said to be produced at around 8% efficiency(*)

I would like to try a well formed PbO2 anode in a divided cell, cooled, using ~50% H2SO4 and a C cathode with saturated (NH4)2SO4 in the anode and high anode current density, later making the potassium salt by double decomposition.(*) During the experiments on electrolytic H2SO4 there seemed to be evidence of persulphate formation on PbO2 even at higher cell temperatures.

(*)Percarbonates (K2C2O6) can be made similarly, using sat K2CO3 solution.

Final notes:

I use terra cotta flower pots as diaphragms. Other permeable membranes can be used; the thinner the better provided they don’t react with any products and are not too fragile.

Current density is rarely critical at the cathode, especially if H2 is evolved. For good metal plating, however, it is critical. Anode current density is usually important if oxidation is concerned. As a general rule, more attention should be paid to current density than any other parameter – a constant current source, however crude, is recommended.

Bath temperature matters with unstable products and cooling is then needed; high temperatures cause a marked decrease in cell resistance. Stirring always helps to avoid anode/cathode build up, especially with insoluble products being oxidized/reduced.

As in many posts on the forum, sprinkle the above with numerous AFAIK, IIRC and IMHOs! Even with the ones I've tried, there's no guarantee of efficiency, but electricity costs are still small.

Enough ideas to be getting on with. Have fun!

Der Alte
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[*] posted on 14-10-2009 at 16:19

Der Alte,
Thank you for a very informative post.
I wish you would consider doing a tutorial for those of us who are not experienced in electrochemical methods. I have done a few simple experiments and have fairly good electrical equipment but I'm just not very conversant with practical details of how to put it all together. I've read some of the books on electrochemistry, but a compendium of practical advice from an expert would benefit me and perhaps a few others. Thanks.
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[*] posted on 15-10-2009 at 05:03


Great post indeed. I did not get time to read it properly so I will be going back to it.

Often wondered would a 'Pb02 Anode' from a car battery do for Sulphuric acid production. I know they have been suggested 1001 times for Chlorate and Perchlorate where they are next to useless. The Lead should stay passivated in Sulphuric solution? and the Lead Dioxide should carry on being useful?

AFAIK keeping current density on the Cathode high in Chlorate or Perchlorate production helps keep reduction (unwanted reactions destroying wanted product) low. Using a small Cathode will raise cell Voltage slightly but we do not care much about that. Manufacturers would care a great deal.

I am going off to purchase some flower pots...


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[*] posted on 15-10-2009 at 14:37

Some nice ideas there DerAlte!

I made several Cu(I) and Cu(II) salts by electrolysis. Generally speaking, electrolysis with salts of "monobasic" acids (e.g. HCl) seem to produce copper(I) whereas salt of "bibasic" acids (e.g. Na2SO4) produce copper(II). CuO can be made most easily by electrolysis of Na2SO4 solution with copper electrodes and subsequent boiling of the precipated copper hydroxide. Electrolysis in NaCl solution seems to produce brown unstable copper(I)hydroxide which can be dissolved in HCl to precipate CuCl upon pouring the solution into cold water.
Also I made as small amount of basic copper(II)carbonate ("malachite") by electrolysis of sodium bicarbonate with copper wire as electrodes. The solution is not terribly conductive thus a high voltage is required to achive acceptable speed. However the overvoltage causes the solution to heat up rapidly. Careful cooling and temperature control is required to avoid decomposition of the copper carbonate into back oxide.

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[*] posted on 15-10-2009 at 16:08

Last time I tried anodic corrosion of copper in Na2CO3, I got a strikingly blue solution: copper forms a carbonate complex which is soluble. Needless to say, this was attended with undesirable copper foam formation at the cathode.

After lots of boiling, azurite I believe will precipitate from this solution. I had left the electrolyte exposed to air, where it formed a blue skin on the surface, probably due to excess CO2 passing out.

I've anodized lead in a fairly strong H2SO4 solution, which resulted in heavy flakes of PbO2 peeling off the anode. Maybe under some conditions (current density, concentration, etc.?) it is self limiting, but that range certainly isn't as wide as one would hope. It's an inefficient way to make bulk PbO2, maybe not bad overall, but it does make excess O2.


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[*] posted on 15-10-2009 at 21:30

Plaese add your electrolytic experiences, like Taoiseach. That will make it a worthwhile thread.

12AX7 wrote

...It's an inefficient way to make bulk PbO2, maybe not bad overall, but it does make excess O2.

Very true! In most cases, the less the better. I have accumulated over 50g of PbO2 as a 'byproduct'! However, it's no mean oxidanr, more of that anon, someday, perhaps. A use for the spare O2 would be good, too. I really ought to make a proper plated PBO2 anode - instead of relying on thin fragile layers.


Most of my apparatus is very crude. A pyrex dish as container or bath, a piece of steel plate or can as cathode, strips, rods or bars as anodes, flower pots as diaphrams. Nothing fancy. If you have a lathe or the tools you can make it more elegant but it will not work that much better! A source of power, preferably of variable voltage, 12v or higher. Most electrolyses work at around the 2 - 3 volt level and a 'constant' current source is a very good thing to have. But you dont need superb regulation, line or load. A 10 to 20% variation hardly matters. I have done many experiments with an unsmoothed 12v battery charger with a lamp as the quasi 'current control' element.

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[*] posted on 1-11-2009 at 09:03

Note 2: nitrates are reduced simultaneously at the cathode to nitrites and ammonia, and possibly also hydroxylamine (NH2OH) so the nitrate case is less efficient.


The reduction of nitrate to nitrite can be accomplished
satisfactorily, and the process is the subject of a recent
I t has been shown (Miiller and "Weber)
tha t in
a divided cell, smooth platinum or copper cathodes reduce
nitrate to nitrite and ammonia, but platinised platinum
gives much ammonia and little nitrite. A spongy copper
or silver cathode was found to give the best results. Wit h
a current density of 0*25 amps, per dm.
and a concentration
of 2*3 grams of sodium nitrate per litre, a current efficiency
of 90 per cent, was obtained. The current efficiency with
an amalgamated copper cathode was found to diminish
when 50 per cent, of the nitrate had been changed.
Considerable care is evidently needed to prevent the forma-
tion of ammonia, since it has been shown by W. H. Easton3 that nitrates may be quantitatively reduced to ammonia by
In the patent referred to above, the cell described is
suitable for the electrolysis of alkali chloride and is of the
bell type, but it is particularly suitable for electrolysing
alkali nitrate. Pure nitric acid is formed at the anode
inside the bell and is removed by distillation, which is
effected by working under reduced pressure and by heating
the bells with superheated steam. The nitrite which is
formed at the cathode is drawn off continuously and
separated outside the cell. The cell itself acts as cathode,
and the anode is of such size as to almost fill the bell and
thus reduce the working space of the electrolyte. High
current density (16 amps, per dm.
), reduced pressure and
high temperature, are favourable to the distillation of a
large amount of concentrated nitric acid.

Source: Page 34 The Manufacturing of Chemicals by Electrolysis.

On the production of Hydroxylamine thru simular means,
Tafel showed that the sulphate is comparatively stable
in the presence of sulphuric acid even at a temperature of
40° C. He obtained the hydrochloride by using hydro-
chloric acid in place of sulphuric acid, and a cathode of
spongy tin gave satisfactory results. The reduction may
be represented by the equation :—
HN0 3 + 3H2 = NH2 0H + 2H2O.
According to the patents of Boehringer and Sohne,
two-compartment cell is employed containing 50 per cent,
sulphuric acid in each compartment. The cathode is of
amalgamated lead, whilst the anode is lead. A 50 per
cent, nitric acid solution is dropped into the cathode com-
partment during the passage of the current, and the tem-
perature kept below 20° by cooling coils. The current
density employed is 60-120 amps, per dm.
According to a French patent,
an anode of platinum
is used with a tin cathode. Sodium nitrate solution is
dropped into the cathode compartment , and the anolyte is
sodium chloride solution. The yield of hydroxylamine is
said to be 60-80 per cent, and chlorine is a by-product.

Source: Page 33 The Manufacturing of chemicals by Electrolysis.

Attachment: 3104731-The-Manufacture-Of-Chemicals-By-Electrolysis[1].pdf (825kB)
This file has been downloaded 1586 times

Very good book and no doubt has its place here in a thread on this very topic. It also goes on to discuss a way to manufacture nitric acid thru the electrolysis of Peat Moss:o. Don't get more Over the counter then going down to your local swamp making nitric acid from the moss then making your nitrates from that followed with a reduction to Hydroxylamine or Nitrite.:D . Hey don't laugh Im sure someone that sees this will try it sooner or later.

[Edited on 1-11-2009 by Sedit]

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[*] posted on 8-11-2009 at 08:36

Excellent reference, Sedit. Has good detail on organic oxidation/reduction plus inorganic. a must read for electrolysers!

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[*] posted on 11-6-2012 at 21:05

An Attempt to produce Nitric Acid by Electrolysis

In the thread ‘Reagent Production by Electrolysis’ supra I wrote:
NITRIC acid is difficult to produce electrolytically from nitrates (except on a Pt anode(*)). All common metal anodes are oxidized and dissolved. Carbon can be tried but also oxidizes - the acid concentration must be kept quite low. Unlike sulphuric acid HNO3 is a strong oxidant even when comparatively dilute. It might be worth a try with a massive deposited PbO2 anode (*). Lead itself is converted to nitrate.

Conclusion: I see no reason to revise that conclusion after another attempt.

It is well known that SS 18/8 or #304 is quite resistant to nitric acid of all strengths due to passivation. I was curious to see what strength nitric acid ,if any, could be produced on an SS anode. The set up used a saturated solution of NaNO3 in the cathode of a divided cell, Fe cathode, SS anode, pure water in the anode cell initially (SS is known to be attacked by chloride ions, hence distilled water).

The cell was run at about 6v, and conduction initiated by dropping a small crystal of NaNO3 into the anode compartment. Current density was low on anode, about 35ma/cm2 and very low at the cathode.

For the first Amp Hour or so it seemed that HNO3 was being produced. Small amounts of gas were emitted at the anode but did not seem to accord with the current drawn. A slight whiff of nitrogen oxides was noted, especially on the anode when withdrawn from solution. The major potential drop was across the cell barrier (Clay pot) but probing the liquid near the anode suggested that a drop at the surface of the anode of about 1.7 volts was occurring. I had hopes that a semiconducting layer was being set up.

However after a further lapse of time the anode solution began to turn reddish. It was obvious that the anode was being eroded. I let the electrolysis run for a few hour and the anode solution became a deep reddish brown color. The solution appeared to be highly acidic. Although the edges of the SS showed erosion, the main surface was still quite bright but on using a magnifier pitting was obvious.

The color did not look right for Chromic acid (CrO3 solution). There may have been some nitric acid also in solution. The cathode cell was a mess of what was obviously Fe(OH)2 (white and gelatinous) plus green blobs of Ni(OH)2 and even a hint of bluish substance (Cr(OH)2 ?). In addition, there was a strong smell of ammonia, which I have noted almost always happens when you electolyse a nitrate, due to reduction at the cathode.

I boiled down the anode solution and finally it gave a deep brick-red powder. At first I though Ferric oxide but the solution was very acidic, assumedly from nitric acid, which was emitted on drying. It then dawned on me that what I had was in all probability ferric dichromate (Fe2(Cr2O7)3); the Cr had been oxidized to CrO4-- and the the Fe to Fe+++. This powder was quite soluble in water and cannot have been hydrated ferric oxide or hydroxide.

An interesting experiment with complicated results, and a total failure in its original purpose!

Regards, Der Alte
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[*] posted on 13-7-2012 at 15:43

A note on alkali hydroxides: stainless steel electrodes on both ends with baking soda as the anolyte and water as the catholyte gives good results.

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[*] posted on 13-7-2012 at 17:32

Hi DerAlte.In your original post you mentioned some experiments oxidising alcohols .Did you enjoy much success with ethanol to acetic acid?

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[*] posted on 6-5-2021 at 18:27

Quote: Originally posted by DerAlte  

With a Cu anode, (D) using saturated NaCl solution as catholyte and NaCl or HCl in the anolyte, instead of the expected CuCl2 passing into anode solution, white CuCl is also precipitated. I suspect the same happens with KBr or KI, giving CuBr or CuI.(*)

Der Alte

After proving this statement experimentally, I am very interested in understanding the mechanism by which this occurs. My research in this regard has been unsuccessful.

When electrolysing a concentrated NaCl solution in a cell divided by a salt bridge, using copper anode and a steel cathode operating at 30V, I obtained an emulsion of CuCl in the anode and only traces of CuCl2.

Why does it happen? My first theory was that "nascent chlorine", that is, monatomic chlorine, was combining directly with the copper at the anode. But it seems to me that the "nascent state" theory is already outdated in chemistry and was abandoned decades ago. In fact, I was so curious about the prevalence of CuCl in this process, since this is the most unstable copper chloride, that I invested my last free afternoon trying to find the mechanistic explanation of why this happens.

Does anybody know?

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[*] posted on 7-5-2021 at 23:44

for metal recovery and plating metals like iron you need a membrane electrolysis cell and the cheap and good PVC substrate PVA polycarboxylate ester membranes can resist up to 35% H2SO4 and are also useful for the process of making H2SO4 from Na2SO4 or any sulfate salt.
PbO2 electrodes are nice and so is Pt but also Ir-Ta MMO can replace PbO2 in some strong acid electrolysis.
Perchlorates need Pt or PbO2 and chlorates and chloro alkali do best with Ir-Ru MMO due to its low chlorine overpotential.
making nitric acid from electrochem directly is garbo!! dont even because it literally will be trash due to decomposition of NO3 at the anode into N2 and O2 even on Pt.
you can however use a copper sacrificial anode with a non porous IEM either commercial or cheap homemade PVC-PVA-Citrate ester membrane.
the copper nitrate can be decomposed as per nurdrage video to make HNO3 and the copper can be dissolved in H2SO4 then plated out to re use the anode.

As for transition metal salts from chlorates and perchlorates this is indeed possible but you still need an IEM to do this effectively which I do have a video for making these cheaply.

[Edited on 8-5-2021 by mysteriusbhoice]
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[*] posted on 8-5-2021 at 10:47

I may have a guess,

Assuming chlorine is the limiting reagent, as it either reacts or bubbles off, there is a excess of copper.

The standard enthalpy of formation of copper(1) chloride is -138KJ/Mol, according to NIST, compared to copper(2) chloride, is -205KJ/Mol. Because of this, it is more favorable for the Cl2 molecule to react and form 2CuCl which a enthaply of formation of -276KJ per mole of Cl2, rather than CuCl2 with a enthalpy of formation of -205KJ per mol of Cl2. If any CuCl2 is formed at the anode, I would guess it would react with copper metal and be reduced to 2CuCl.
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