Sciencemadness Discussion Board - Copper Citrate (and other transition metal citrates)

Copper Citrate (and other transition metal citrates)

Romix - 1-7-2015 at 11:20

Pure copper anode and cathode, electrolysed through sodium carbonate.
Blue hydroxide formed on anode. Had to change polarity every 10 minuts, cos it sticks to electrode, and blocks conductivity.

This hydroxide I dissolved in citric acid, produced alot of bubles. I got told that what I call hydroxide is basic copper carbonate. And bubles produced are carbon dioxide.

That's how I balanced equation for this reaction 2C6H8O7 + 3Cu(OH)2 = Cu3(C 6H5O7)2 + 6H2O
There should'nt be any bubles.

[Edited on 1-7-2015 by Romix]

[Edited on 1-7-2015 by Romix]

[Edited on 1-7-2015 by Romix]

Edit by Texium: changed title with thread merge, for clarity

[Edited on 9-14-2023 by Texium]

papaya - 1-7-2015 at 13:12

You produced malachite, that's why it's bubbling, however the exact chemical structure of that deposit might be different and cannot be reliably identified (how much hydroxide vs basic carbonate is formed, how hydrated is it, OH- anions sometimes serve as ligands for transition metals, sometimes bridge type ligand, etc,etc..). For such type of "compounds" I have one description only - sludge!

blogfast25 - 1-7-2015 at 13:31

Quote: Originally posted by papaya

You produced malachite, that's why it's bubbling, however the exact chemical structure of that deposit might be different and cannot be reliably identified (how much hydroxide vs basic carbonate is formed, how hydrated is it, OH- anions sometimes serve as ligands for transition metals, sometimes bridge type ligand, etc,etc..). For such type of "compounds" I have one description only - sludge!

No, no. What you write here is factually very incorrect. The composition of Malachite is well defined: equimolar Cu(OH)2 + CuCO3 = Cu2CO3(OH)2.

https://en.wikipedia.org/wiki/Malachite: scroll down to half way down for a ball and stick model.

There's NO reason to believe that electrolytic basic copper carbonate would have a different structure or have a different hydroxide/carbonate ratio. Water isn't chemically bound to it, so that's different.

In Romix's process, Cu2+ is formed through electrolytic oxidation. Then:

2 Cu2+(aq) + 2 CO32-(aq) + H2O(l) === > Cu2CO3(OH)2(s) + CO2(g)

That compound is definitely NOT 'sludge'!

[Edited on 1-7-2015 by blogfast25]

papaya - 1-7-2015 at 14:09

I agree that the malachite is well defined, but it's still possible that under not very well controlled conditions what will form is not only malachite. Look at your equation
2 Cu2+(aq) + 2 CO32-(aq) + H2O(l) === > Cu2CO3(OH)2(s) + CO2(g)

The release of CO2 means that if you use small volume of carbonate solution it will probably lower the pH over the time and go more to bicarbonate side, what happens then I don't know. But if you have information that this process is really reliable for malachite production - I cannot object.

blogfast25 - 1-7-2015 at 15:03

Quote: Originally posted by papaya

The release of CO2 means that if you use small volume of carbonate solution it will probably lower the pH over the time and go more to bicarbonate side, what happens then I don't know. But if you have information that this process is really reliable for malachite production - I cannot object.

Sigh... Malachite is formed BOTH by carbonate AND bicarbonate solutions. Only pure hydroxide forms Cu(OH)2.

2 Cu2+ + 4 HCO3- === > Cu2CO3(OH)2 + 3 CO2 + H2O

So with bicarbonate you're simply wasting more CO2.

The only way you might run into Cu(OH)2 with Romix's process is by running out of carbonate or bicarbonate and then you're simply making malachite first, then copper hydroxide.

There's an awful lot of misinformation on the web re. copper carbonate, carbonates and bicarbonate, spread by amateurs that make salient claims without ever having measured a g-ddamn thing!

Also, released CO2 cannot lower the pH of a carbonate solution. Small or large volume has nothing to do with anything: we're talking concentration here.

[Edited on 1-7-2015 by blogfast25]

Romix - 1-7-2015 at 15:13

I don't understand how basic copper carbonate formed.

blogfast25 - 1-7-2015 at 15:17

Quote: Originally posted by Romix

I don't understand how basic copper carbonate formed.

If you can't understand the explanation offered above then I can't help you, sorry.

blogfast25 - 1-7-2015 at 18:10

Quote: Originally posted by Romix

That's how I balanced equation for this reaction 2C6H8O7 + 3Cu(OH)2 = Cu3(C 6H5O7)2 + 6H2O

I also think you prepared dicopper citrate, not tricopper citrate (as you claim). I could be wrong on that.

[Edited on 2-7-2015 by blogfast25]

violet sin - 1-7-2015 at 18:29

Romix: here ya go, if you don't see where the hydroxide ion is coming from, to quote blogfast25:
----------
n Romix's process, Cu2+is formed through electrolytic oxidation. Then:
2 Cu2+(aq) + 2 CO₃2-(aq) + H₂O(l) === > Cu₂CO₃(OH)₂(s) + CO₂(g)
----------
one of the CO₃2-(aq) has to release an oxygen to become CO₂(g) , and water H₂O(l) gains that to become 2 OH-(aq)
CO₃2-(aq) + H₂O(l) -> CO₂(g) + 2 OH-(aq)
the charge is maintained, but i'm a bit rusty on explaining *exactly how/why/where that happens,

I have also heard people say (on youtube so grain of salt here) that the Na+(aq) participates by traveling to the electrode, receiving an e- -> Na0 and instantaneously 2 [ Na0 + H₂O(l) ] -> 2 Na+(aq) + 2 OH-(aq) + H₂(g)
^^^( open for criticism)^^^

man, it is a bit tedious to do all the sub/sup by hand...

Romix - 1-7-2015 at 18:52

Oxygen on anode, hydrogen on cathode, copper oxidized to Cu + 2 , not going far drops to the bottom of electrode, and stick to it after a while blocking conductivity.

[Edited on 2-7-2015 by Romix]

violet sin - 1-7-2015 at 20:43

https://en.m.wikipedia.org/wiki/Basic_copper_carbonate
Basic copper carbonate is insoluble,.. so If you have hydroxide and carbonate anions roaming free, and they happen across a swarm of copper cations... Theyz get married and settle down.

blogfast25 - 2-7-2015 at 04:23

Quote: Originally posted by violet sin

https://en.m.wikipedia.org/wiki/Basic_copper_carbonate
Basic copper carbonate is insoluble,.. so If you have hydroxide and carbonate anions roaming free, and they happen across a swarm of copper cations... Theyz get married and settle down.

This is kind of the key point: Cu2CO3(OH)2 is more insoluble than Cu(OH)2 or CuCO3, that's why it precipitates preferentially, if carbonate/bicarbonate is present.

sasan - 2-7-2015 at 08:04

Hi,there is another method using sodium citrate and copper sulfate with 1 : 1.5 molar ratio.make the solutions and treat with each other.leave it for a night and filter the precipitate.
http://www.rsc.org/learn-chemistry/resource/download/res0000...

blogfast25 - 2-7-2015 at 09:14

Quote: Originally posted by sasan

Hi,there is another method using sodium citrate and copper sulfate with 1 : 1.5 molar ratio.make the solutions and treat with each other.leave it for a night and filter the precipitate.
http://www.rsc.org/learn-chemistry/resource/download/res0000...

Yep.

I'm not convinced that citric acid is strong enough to fully neutralise basic copper carbonate.

The third deprotonation of citric acid (H3Ct) has a pKa3 of only 6.4, that's extremely weak and copper basic carbonate is NOT a strong alkali.

To be clear, the third dissociation:

HCt2- + H2O <=== > H3O+ + Ct3-

... is extremely weak.

Did Romix actually obtain a precipitate?

[Edited on 2-7-2015 by blogfast25]

papaya - 3-7-2015 at 00:26

Citric acid is a very strong complexant, it doesn't need to be a strong acid to dissolve carbonate, but of course you know it blogfast!

violet sin - 3-7-2015 at 05:31

tonight I checked the (supposed) copper hydroxide I made a few years back, still light sky blue in it's mason jar after all this time. it did not bubble the least with strong HCl( picture below just moments after addition). pool/concrete etchant strength from the hardware store. I got the how to from youtube as copper electrodes in a dilute(ish) aqueous epsom salt bath, with a hacked ATX computer PSU. the sample instantly turned a beautiful clear emerald green when the HCl was poured in, dissolving the clumps rather fast.

after I cleaned the test tube out, pleased that it didn't fizz, I also put some of the copper hydroxide in an iron citrate solution I had just made. [ rust and a little steel wool in excess citric acid, all dissolved; made a clear darker yellow sol pictured below, top two]. the gloppy blue copper hydroxide paste( I had kept it wet) was shook in a few ml of the iron citrate. after the cloudiness settled out in a few min, the solution began turning rather green( pictured below, bottom 2). I honestly don't know how much free citric acid was left in the iron citrate sol, or how much rust and steel wool I had in there. I was just hoping to make some crystals later.

so, no I didn't make copper citrate directly, but I did test the copper hydroxide I made(with HCl) and then get the iron citrate I had already made, to take up some copper. not a bubble in either case.

....................................................

....................................................

blogfast25 - 3-7-2015 at 06:09

Quote: Originally posted by papaya

Citric acid is a very strong complexant, it doesn't need to be a strong acid to dissolve carbonate, but of course you know it blogfast!

I hadn't thought of that, papaya, no. Ta.

blogfast25 - 3-7-2015 at 06:17

Violet sin:

The Fe citrate complex (FeCt) is likely stronger than Cu3Ct2. Thus no ligand transfer occurs:

FeCt + CuCl2 === > Fe chloride + Cu citrate doesn't happen.

Try neutralising some NaOH with citric acid, then add this to your CuCl2 solution, that's more or less what 'sasan' linked to above. Na+ cannot be complexed by citrate ions, so you have free Ct3- anions in there.

[Edited on 3-7-2015 by blogfast25]

blogfast25 - 3-7-2015 at 08:45

For the benefit of ‘violet sin’ (his request via U2U):

The equilibria that rule a solution that contains a dissociated copper(II) salt a soluble carbonate are:

1) Re-protonation of carbonate anions:

CO3(2-) + H2O < ==== > HCO3- + OH-

Leans much to the left but causes the solution to be alkaline.

2) Auto-dissociation of water:

2 H2O < === > H3O+ + OH-

Leans very much to the left.

3) Re-protonation of bicarbonate anions:

HCO3- + H3O+ === > H2CO3

Leans much to the right.

4) Decomposition of carbonic acid:

H2CO3 < === > H2O + CO2(g)

Leans much to the right.

5) Solubility product of basic copper carbonate:

Cu2CO3(OH)2 < ==== > 2 Cu2+ + CO3(2-) + 2 OH-

Leans much to the left.

One can set up the relevant equilibrium equations (including equilibrium constants, all are known), including mass balance and neutrality balance and solve this system of simultaneous non-linear equations for the concentrations of all species and would conclude that:

2 Cu2+ + 2 CO3(2-) + H2O === > Cu2CO3(OH)2(s) + CO2(g)

… proceeds as observed.

More simply put, the formation of an insoluble reaction product and a gaseous one, ensures that the change in Gibbs Free Energy ΔG (left to right for the overall reaction) is negative, i.e. favourable.

violet sin - 3-7-2015 at 13:57

-The mix was copper hydroxide, iron citrate with an excess of citric acid. And it turned green. Either the excess citric acid in solution dissolved more of the Cu(OH)₂ overnight, or it settled more compactly. I'll add some more citric acid to see if it will eat all the hydroxide. Oh ya, there is no brown ppt in there. Indicating the copper hydroxide didn't displace the iron, both are in sol.

-I had the iron citrate sitting around b/c I wanted some crystals of just that. It wasn't purpose made for this experiment, just convienient.

-The copper chloride was dumped in my copper chloride e-waste jug. And did not participate here at all. Just for the carbonate test.
----------
all the copper hydroxide dissolved leaving a clear green solution.

[Edited on 4-7-2015 by violet sin]

AJKOER - 5-7-2015 at 08:32

I have prepare Copper citrate fairly quickly by simply combining Copper metal, Citrate acid, dilute H2O2 and a touch of sea salt (or plain salt) to act as a good electrolyte for the galvanic cell reaction (think/research metal-air battery where the H2O2 replaces O2, and with an acidic electrolyte, the strength of the conjugate base, dihydrogencitrate, is involved).

I usually jump start the reaction by heating the mixture in a microwave, at which point, the reaction actually appears to be quite vigorous. In minutes, my experience is that the solution looks like the above pictures.

No external current or electrodes required or even continuous heating, for that matter. You may have to add more H2O2 and reheat periodically until the Citrate acid, is completely consumed, however.

Note, the final product will have some NaCl impurity, which if is an issue for an intended application, could be addressed through recrystallization. Also, ones copper source may be an alloy, resulting in yet another source of possible impurities. Even very high purity electrical Copper wiring, for example, may contain some Cadium (see discussion at http://www.copper.org/resources/properties/microstructure/ca... ).

[Edited on 5-7-2015 by AJKOER]

mayko - 31-8-2015 at 17:57

Here are some notes from my explorations of copper citrate. The preparation came from the RSC document in the Sources section:

Experimental and Observations
32.11 g (0.157 mol; 1 eq) citric acid monohydrate was added to a 250 mL beaker containing 150 mL distilled water, and the mixture was heated to dissolve. The pH of the solution, measured with multi-indicator test strips, was ~1.5. 22 g of sodium hydroxide were added (0.55 mol, or 3.7 eq, assuming 0% moisture in the hydroxide), bringing the pH to ~12. The liquid was poured into an evaporating dish and heated until bubbling gently and slightly syrupy. As it cooled, the liquid began to form a sort of 'pudding skin' on the surface. In the unheated lab, the mixture froze solid overnight; on warming, it remained a solid cake of white crystalline solid. This was broken up and added to a 250 mL beaker containing 50 mL of 200 proof ethanol. The slush was stirred, filtered and rinsed with another 50 mL dry ethanol. Once the solid had drained thoroughly, it was crushed finely and stirred into 100 mL boiling dry ethanol. After several minutes, it was filtered and laid out to dry. The white, fluffy crystalline solid massed at 46.24 g (0.157 mol; 1 eq; i.e., nearly qualitative yield.) A solution of the salt was measured at pH~10, compared to a literature value of 8, suggesting residual hydroxide contamination. Adding copper sulfate to the solution did not yield precipitate; the mixture appeared to darken with small dilution.

In an erlenmeyer flask, 5.91g sodium citrate (0.02 mol, 1eq) and 7.54g copper sulfate pentahydrate ( 0.03mol. 1 eq) were dissolved in steaming distilled water, creating a definite darkening of the copper blue to a deep teal. On standing, a scummy deposit, not a grained precipitate, formed. Upon heating, the solution precipitated a fine turquoise powder. This was washed with alcohol, dried, and massed at 3.39 g . Taking at face value the hydration state in Merck of 2.5 and the hydrate molar weight of 613 g/mol, this is ~0.0055 mol, ie 55% yield. (There were mechanical losses when the flask burped in the microwave.)

To confirm the hydration state, 0.50 g copper citrate were placed in a test tube and heated in a 50 mL beaker containing mineral oil (dehydration occurs at 100C, says Merck). Moisture was seen leaving the tube, and the solid changed from a sea-foam green to a sky blue. The pale dehydrated salt was massed at 0.46 g (0.81mmol). This suggests 0.04g water (2.2mmol) were driven off, giving a ratio of ~2.75 water molecules per copper citrate molecule, agreeing with the literature value.

A small amount was placed in a test tube and heated while continuously flushed with carbon dioxide. Granular metallic copper resulted, though no pyrophoricity was observed.

RSC provides an alternative procedure, based upon copper acetate and citric acid, rather than copper sulfate and trisodium citrate. In a 50 mL flask, 1.50g copper acetate (8.2mmol) and 1.09g (5.1mmol) were dissolved in ~20 mL distilled water. On standing, a small amount of green crust formed. When the liquid was poured into a new flask and heated to boiling, it again became turbid green and precipitated. (Below, left) Magnesium citrate, available in solution at the pharmacy as a laxative, will also precipitate when heated with copper sulfate. (Below, right)

Merck claims that copper citrate is soluble in hot solutions of alkali citrates; this was attempted with hopes of recrystallization. In 25 mL of boiling distilled water were dissolved 3.4 g trisodium citrate, and 1.6 g copper citrate was added. The mixture was at first turbid and murky, then became a dark, clear blue solution. This was filtered hot into a new 50 mL flask cooled to room temperature, then chilled. After two days, no crystals were observed; the flask was heated down to ~5mL to drive off water. This was chilled another day, but only became more syrupy. When a drop was placed in a test tube of pure water, or a 50/50 mixture of water and ethanol, the liquid became a clear blue solution; in 100% ethanol, the drop settled on the bottom like a bead and stuck to the glass. When ~1mL 100% ethanol was added to the flask containing the syrup, it formed a bilayer (!) which would reappear even when mixed with vigorous shaking (!!). Over several months, the liquid solidified into a pale blue crust. The solubility of copper citrate in concentrated ammonia was also confirmed, forming a deep tetraamine-copper blue solution.

Data & Materials

Citric acid C6H8O7 (mw=192.12 g/mol). Exists as a monohydrate (mw = 210.14g/mol); LD Carlson Co. product was used, via 5th Season, a local homebrew store.
Trisodium citrate Na3C6H5O7 (mw = 258.07 g/mol) Exists as a dihydrate (mw = 294.1 g/mol). Soluble in water, insoluble in alcohol
Copper (II) Citrate Cu3(C6H5O7)2 ; Cu3C12H10O14. mw = 568.85 g/mol . Hydration state claimed by Merck is 2.5 (mw = 613 g/mol)
Sodium Hydroxide NaOH (mw = 40g/mol) Comstar “99% Pure Lye“ was used. Slightly soluble in alcohol.
Copper acetate (mw=181.6 g/mol)
Copper sulfate pentahydrate. (mw = 250 g/mol)
Magnesium Citrate. Store-brand lemon-flavored laxative solution was used.

Discussion

This project grows out of an interest in pyrophoric copper, which can be prepared by heating in vacuum according to Gorrie, Kopf, & Toby (1967) (h/t ScienceMadness user Pok). Lacking a vacuum pump, this could not be replicated using a CO2 atmosphere, although metallic copper did appear to result from thermal decomposition. Although the original purpose of this investigation (pyrophoric copper) was not achieved, several interesting observations were made. Some of these suggest directions for future exploration.

The first is misinformation and its propagation in the literature. Neither protocol given in RSC would work satisfactorily without additional heating. More severe is an error present in both the Merck Index [8th ed] and the CRC [63rd ed], listing the formula for copper (II) citrate as “Cu2C6H4O7*2.5H2O”, with a molar weight of “63.220”. Although casual consideration of the citric acid acid molecule suggests this is unlikely (it would require the deprotonation of a hydroxyl group in addition to three carboxylic acid groups), this formula has spread to internet databases and peer-reviewed literature The CRC still lists the incorrect formula at the time of writing [96th Ed, 2015-2016]

When prepared, the sodium citrate seemed to thicken and gel, rather than crystallize, then solidified into a block. When copper citrate was dissolved in a sodium citrate solution and evaporated, it formed a syrupy fluid which slowly solidified. Perhaps in both cases, the salt(s) dissolved in the water of hydration upon heating. Could this explain the behavior of the copper solution, if the salts were dissolved in too much water to form an anhydrous solid, but too little to form the proper hydrate, forcing it into a fluid state until sufficient moisture could be absorbed from the atmosphere? I have heard similar reports from attempted crystallizations of laxative magnesium citrate solution.

The darkening of the reaction mixture in spite of a dilution of the copper ion, would be an interesting area for spectrophotometric investigation.

References

Preparation of copper(II) citrate. RSC Student worksheet, http://www.rsc.org/learn-chemistry/resource/download/res0000...

CRC Handbook of Chemistry and Physics

Merck Index, 8th Edition

Gorrie, T., Kopf, P., & Toby, S. (1967). Kinetics of the reaction of some pyrophoric metals with oxygen. The Journal of Physical Chemistry, 1478(4), 3842–3845. Retrieved from http://pubs.acs.org/doi/pdf/10.1021/j100871a019

mayko - 20-10-2016 at 17:27

Quote: Originally posted by mayko

Magnesium citrate, available in solution at the pharmacy as a laxative, will also precipitate when heated with copper sulfate.

I realized the other day that there are actually two distinct magnesium citrate salts:
*there is a less soluble trimagnesium citrate, with the formula Mg₃Cit, and
*the sort that is used in this laxative, which is referred to as magnesium citrate but might be better called magnesium hydrogen citrate, MgHCit.

It occurred to me that I'd never confirmed that the copper salt I prepared from the laxative solution was the 3:2 salt, rather than a 1:1 hydrogen salt like CuHCit.

To find out, I heated down 30 mL of the laxative (at ~1 fluid oz; at 1.745 g/flOz and 214 g/mol, that's ~0.008 mol) to drive off some water, and dissolved ~3 g copper sulfate pentahydrate (approx. stoichiometric) in a minimum of water. a bit of each solution was touched to some universal pH paper, with no color change except a bit of bluing from the copper sulfate. The two solutions were then mixed and heated until precipitation occurred. The solution was checked again with indicator paper, which now turned orange-red.

If the product is the 3:2 salt, the citric acid must lose both a magnesium cation and a proton. If it's the 1:1 hydrogen salt, it will only add magnesium to solution. The change in pH from approximate neutrality of the reactant solutions to quite acidic in the product supports the 3:2 salt.

Amos - 12-12-2016 at 14:23

I love how mayko swiftly ended all the bickering and conjecture above and nobody said anything about it. Mayko, thank you for such an exhaustive report!

I've had two perplexing vials of both the anhydrous and hydrated forms of this copper citrate complex sitting on my lab windowsill for what seems like an eternity that I inadvertently produced using a powdered drink mix... yeah. I was completely stumped at the time as to what they could've been and before long I had forgotten how I obtained them in the first place, until seeing your pictures in the wiki reminded me of that afternoon where a less-learned me poured lemon-flavored Kool-Aid powder into some copper salt solution in a half-witted attempt to make copper(I) oxide. Anyway, thanks again!

[Edited on 12-12-2016 by Amos]

mayko - 10-2-2018 at 23:22

Quote: Originally posted by mayko

The 2.5 hydration multiplicity kept sticking out in my mind and I'm starting to doubt it. My earlier gravimetric measurement of the hydration state was based on a relatively small mass (~half gram) with a low precision (0.01g) scale; additionally, the water molecules are much less massive than the copper citrate (by a factor of ~30), meaning the difference in mass between the hydrated and dehydrated samples was miniscule (0.04g). Assuming this difference is accurate to +/- 0.01g, this experiment could actually support a range of hydration multiplicities from 2 to 3.5.

I've done a few experiments with larger starting masses and they have generally indicated a multiplicity of 2. For example:

9.53g Cu3Cit2 * nH2O were heated to dehydrate the remaining Cu3Cit2 massed at 8.96 g (ie, 15.8 mmol). The difference of 0.57 g (ie, 31.6 mmol water) gives a hydration multiplicity of n=2.01.

The main source of error, I'd expect, is from mechanical losses; these would tend to bias the calculated multiplicity upwards (ie, greater than 2.5). Contamination or incomplete dehydration might biase the value downwards (ie, below 2.5) but I've come up with a value pretty close to 2 several times now.

Quote:

This project grows out of an interest in pyrophoric copper, which can be prepared by heating in vacuum according to Gorrie, Kopf, & Toby (1967) (h/t ScienceMadness user Pok). Lacking a vacuum pump, this could not be replicated using a CO2 atmosphere, although metallic copper did appear to result from thermal decomposition. Although the original purpose of this investigation (pyrophoric copper) was not achieved,

This sort of depends on your definition of pyrophoricity; I still haven't observed the powder catch fire when poured out cold, but when poured h

Quote:

The solubility of copper citrate in concentrated ammonia was also confirmed, forming a deep tetraamine-copper blue solution.

I was able to isolate this as a dark blue solid two different ways. In the first, I would dissolve copper citrate in ammonia solution, and then add dry alcohol (91%IPA), forming a bilayer similar to the one pictured above. This was mixed, left to stand, and the alcohol replaced. Eventually, the complex formed a solid crust, which was further dried by grinding under dry ethanol:

IMG_20170303_cuCit_Crust.scaled.jpg - 332kB

This is problematic for a lot of reasons: much too lossy for gravimetry; takes a lot of alcohol; time consuming; raw product is quite tough and difficult to dislodge. So, I tried another method, in which copper citrate was placed in a large test tube, covered in dry alcohol, and ammonia gas bubbled through. The suspension becomes flocculent and darkens; on standing, a dark blue solid remains.

IMG_20170413_cucit_aminator.scaled.jpg - 373kB

The experimental apparatus

IMG_20170416_cucit_amination_early.scaled.jpg - 274kB

IMG_20170416_cucit_amination_late.scaled.jpg - 275kB

The reaction mixture early (left) and late (right) in the process.

IMG_20180205_cuCit_amine_compare.scaled.jpg - 394kB

Top to bottom: Hydrated, dehydrated, and ammonia complex.

I've tried a few times to measure the multiplicity for the ammonia complex, and haven't been able to get much of anywhere. I think a number of factors are working against gravimetry as a tool here, but it's not obvious what else is available.

Junk gravimetry results:

Code:


  	Copper Citrate, anhydrous equivalent	  	  	Calculated amination multiplicity, given:	 
Copper Citrate mass (g), hydrated	hydration multiplicity = 2	hydration multiplicity= 2.5	Aminated Copper Citrate	hydration multiplicity = 2	hydration multiplicity= 2.5
1.05	0.99	0.97	0.9	-2.96	-2.5
1.01	0.95	0.94	0.98	1.06	1.57
5.01	4.71	4.64	4.64	-0.51	0

Anhydrous starting material:

Code:


Copper citrate mass, anh	Aminated Copper Citrate	calculated amination multiplicity			
1.04	0.96	-0.58

Preparing soluble metal citrates

KoiosPhoebus - 19-8-2023 at 01:39

Citric acid is a widely available, weak organic acid capable of chelating most multivalent metals. This occurs due to its three carboxyl groups which are negatively-charged when deprotonated (diagram from Zabistak et al, 2018):

In addition to being widely available, citric acid is also relatively safe, biodegradable, and one of the stronger naturally-occurring chelators (Table of four organic acid complex stability constants, Resource on stability constants of chelators used in food, Stability constants of some other organic acids). Metal citrates, however, are not nearly as common as, say, metal gluconates, due to their insolubility. One reason is that citric acid is a triprotic acid, meaning that when a divalent metal hydroxide is reacted with citric acid, a polymeric structure is formed with the formula Metal₃(Citrate)₂ (see row a in the image above). This structure is typically insoluble and precipitates out of solution - which means that adding stoichiometrically equal amounts of the divalent metal and citric acid doesn't work:

3 Metal(OH)₂ + 3 HOC(COOH)(CH₂COOH)₂ = Metal₃(HOC(COO)(CH₂COO)₂)₂ ↓ + HOC(COOH)(CH₂COOH)₂ + 6 H₂O

(The equation above is heavily simplified; in reality there would be a mix of citric acid, dihydrogen or hydrogen citrate ions and some of the metal would remain in solution as either free ions or complexed)

I was interested in using a chelated form of copper as an algaecide and as a soluble fungicide for my plants. So I thought, why not try adding an ammonia molecule to counter-balance the charge? i.e.:

Cu(OH)₂ + NH₃ + HOC(COOH)(CH₂COOH)₂ = CuNH₄(HOC(COO)(CH₂COO)₂)

As monovalent cations are only weakly chelated by most chelators, in theory, when the copper-ammonium citrate is added to water, the ammonium ion should dissociate and result in a negatively-charged complex which readily dissolves. I got the idea from agricultural metal chelates, where they have extra sodium to counterbalance the 4- charge of EDTA (e.g. "iron EDTA chelate" is usually FeNaEDTA).

In an attempt to prepare copper ammonium citrate, I dissolved 9.71 g of anhydrous citric acid in about 30 mL of distilled water and used a magnetic stirring disc to dissolve all the citric acid.

Then, I added about 4.63 mL of 20% aqueous ammonia (having checked the density before use) and let it react with the citric acid. The solution warmed up as the acid-base reaction occurred.

After about 5 minutes, I added about 5.03 g of copper hydroxide I prepared from copper sulphate using the method mentioned in NileRed's Making copper hydroxide video. The copper hydroxide dissolved and formed a dark, blue-green complex:

Unfortunately, after some time, a precipitate formed on the sides and bottom of the beaker.

Since a significant amount of the copper appeared to remain in solution, my suspicion was that the first reaction I mentioned had occurred, and tribasic copper citrate (Cu₃(HOC(COO)(CH₂COO)₂)₂) had precipitated. This would have resulted in an excess of citric acid in the solution, which lowered the pH and kept the remaining citrates somewhat protonated.

In response, I added 4.63 mL of the 20% ammonia solution, twice (in a later run, I added it just once and the precipitate eventually dissolved), and transferred it to a crystallising dish. The solution became a deep, dark blue, suggesting that some of the ammonia had complexed the copper ion.

I then boiled the solution down on a hotplate. In the run where I added 2x 4.63 mL of 20% ammonia, ammonia volatilised from the solution (causing me to move it outside), while in the run where I added only 1x 4.63 mL of 20% ammonia, practically no ammonia volatilised. The presence of ammonia was tested using a wet pH paper as well as checking the pH of a nearby dish of citric acid.

Once the solution had become thick, I took it off the heat and let it cool down. When it was at room temperature, I added denatured absolute ethanol, which initially resulted in a liquid bilayer of the copper/ammonia/citrate solution at the bottom. However, with stirring and mixing using a spatula, the copper-ammonium citrate crystallised as a blue-green solid. Interestingly, in my first run, not all of it crystallised, with a small portion remaining as a thick blue goo despite many attempts to scratch at it.

I then filtered off the precipitate, and dried it on a hotplate at around 50 deg C. Interestingly a small amount of ammonia seems to have volatilised at this point, but I'm not sure if it was from ammonium in the crystal structure of the copper-ammonium citrate, or from excess ammonia which had dissolved into the ethanol etc.

The final product was about 15.45 g of a copper ammonium citrate which looked like a slightly lighter version of copper hydroxide.

The product is water-soluble, with no signs of any precipitation after about 5 minutes. The colour of the solution reminds me of copper gluconate; which makes sense as both citric acid and gluconic acid coordinate to metals via -COO and -OH groups.

Additionally, at least some of the ammonia present in the copper-ammonium citrate is present as the protonated ammonium form. When I added some potassium hydroxide to the solution above, it became a much darker blue typical of the copper-ammonia complex. The stoichiometry of the product is unclear due to the losses of ammonia during the precipitation and drying process, but I intend to estimate the amount of copper in the compound by precipitating it as cupric oxide.

In the future, I might do the ammonium citrates of manganese, nickel and cobalt. The only challenge is that I'm not sure if a metal carbonate will work as well as the metal hydroxide. An initial attempt using ammonium bicarbonate instead of aqueous ammonia to prepare the initial ammonium dihydrogen citrate resulted in green-blue precipitate that didn't dissolve even when aqueous ammonia was added. It might be the case that some bicarbonate ions remained in solution and precipitated a copper carbonate compound. Unfortunately, some metal hydroxides are very difficult to prepare dry/pure due to oxidation (Mn(OH)₂ is an example of this, not sure about nickel or cobalt) so I'll have to try and use the carbonate for such metals.

[Edited on 19-8-2023 by KoiosPhoebus]

[Edited on 19-8-2023 by KoiosPhoebus]

[Edited on 20-8-2023 by KoiosPhoebus]

Admagistr - 19-8-2023 at 12:44

I recently opened this thread on the Forum, it would be a good idea to merge it with yours, so that it is not here twice...It is then unnecessary and uncluttered, when one topic on the Forum is duplicated...

Synthesis of ferrous ammonium citrate

KoiosPhoebus - 13-9-2023 at 22:30

The next soluble metal citrate I attempted to prepare was iron(II) ammonium citrate. Upon some research, I found the compound sodium ferrous citrate, so I decided to use a 1:2 ratio of iron to citrate, similar to sodium ferrous citrate (plus 4 ammonium cations to balance the charge on the complex). Ferrous citrate (Fe₃(Citrate)₂) is prone to oxidation in air, so I thought the slight acidity contributed by the ammonium ions would help keep it air-stable (similar to ferrous ammonium sulphate).

To minimise oxidation, I boiled about 200 mL of water to remove as much oxygen as possible. When the water cooled to about 90 deg C, I added 19.4 g of citric acid (~0.1 mol) and stirred until it fully dissolved. The magnetic stir bar was then removed and 2.85 g of fine mesh iron powder (~ 0.05 mol) was added. A large amount of bubbling occurred as the metallic iron reduced the H⁺ to H₂. After some time to allow the hydrogen gas to displace oxygen from the flask, I sealed the flask with a cap to prevent oxygen from entering and oxidising the iron(II) ions.

Fe + 2 HOC(COOH)(CH₂COOH)₂ = Fe(HOC(COO)(CH₂COOH)₂)₂ + H₂

I left the solution to react for about a day; when I came back, it had a dirty-green-yellow appearance:

There was some residual black powder at the bottom of the flask; however, it did not respond to any magnets I placed near it and so it's probably carbon impurities and I assumed the reaction completed.

I then added 18.5 mL of 20% aqueous ammonia (~0.2 mol) to deprotonate the dihydrogen citrate and form ferrous ammonium citrate:

Fe(HOC(COO)(CH₂COOH)₂)₂ + 4 NH₃ = Fe(NH₄)₄(HOC(COO)(CH₂COO)₂)₂

This caused the solution to change from yellow-green to a dark yellowish brown:

This surprised me, and I was worried that the iron had been oxidised. However, when I pipetted a small amount of the solution into a concentrated solution of sodium hydroxide, the precipitate formed was green (before gradually becoming a red-orange), which suggests to me that the iron in the complex is still in its +2 oxidation state:

The solution retained its colour even when the impurities were filtered off with a double layer of coffee filters.

After some reading on other chelators which coordinate via hydroxyl and carboxyl groups to iron, I think this might actually be normal (google ferrous fumarate powder, or ferrous gluconate powder for examples). The usual yellow-green of iron(II)/citric acid solutions might be due to the citric acid being partially protonated and therefore coordinating differently.

I then placed the solution into a crystallising dish and began heating it to evaporate off the water. I added a small amount of iron powder to the solution to avoid any oxidation of the iron(II) complex; reasoning that I could remove most of it with a magnet after. When the solution was boiled down to about a fifth of its initial volume, I added ethanol to precipitate the ferrous ammonium complex.

When ethanol was added and the liquid was mixed with a spatula, a chocolate-brown flocculent solid initially formed. However, after some time, it changed colour to a dark yellow-green:

The remaining ethanol/water solution was a bright yellow-green, suggesting to me that there was still some iron citrate dissolved in there:

I then dried the yellow-green precipitate and obtained a bright yellow-green solid. The solid was very hygroscopic - even when placed in a sealed ziploc bag with a silica gel desiccant, it still caked:

(fresh iron dihydrogen citrate solution on the left)

However, I'm not confident that this is ferrous tetra-ammonium bis-citrate. While I'm confident that the iron remains in its +2 oxidation state, I suspect that some of the ammonium may have somehow reverted to ammonia and was removed by the ethanol.

This is what the product looks like, dissolved in water:

Adding excess citric acid results in a brighter yellow-green solution:

(lighting isn't exactly the same unfortunately, but I can attest to a colour change)

On the other hand, adding tri-potassium citrate changes the colour to brown, reminiscent of other ferrous complexes:

Fully-deprotonated citrate ions are stronger chelators than protonated citrate (data for Fe(III): https://www.researchgate.net/figure/Stability-Constants-for-... data for Cu(II): https://www.researchgate.net/figure/Stability-constants-of-c...) and so they can displace protonated citrate species from the complex. I suspect this might be what caused the colour change.

There's also the possibility that the iron in the complex may actually have oxidised at some point, but exposure to light resulted in a redox reaction between oxidised iron(III) and the citrate, resulting in iron(II) acetonedicarboxylate. My current plan is to repeat the procedure but with the more basic potassium instead of ammonium, and see what happens.

Synthesis of tetrapotassium ferrous citrate

KoiosPhoebus - 24-9-2023 at 02:27

Similar to my above attempt at preparing ferrous ammonium citrate, I first boiled 200 mL of water, waited for it to cool down to 90 deg C, and then dissolved 19.4 g of citric acid (~0.1 mol) into the solution followed by the addition of 2.85 g of iron powder (~0.05 mol).

Once the effervescence was complete, and any remaining particles no longer responded to a magnetic field, I dissolved 12.47 g of 90% potassium hydroxide into 50 mL of water and added that into the flask. The reaction being:

Fe + 2 HOC(COOH)(CH₂COOH)₂ = Fe(HOC(COO)(CH₂COOH)₂)₂ + H₂ ↑

Fe(HOC(COO)(CH₂COOH)₂)₂ + 4 KOH = K₄Fe(HOC(COO)(CH₂COO)₂)₂ + 4 H₂O

(Note that it's preferable to pre-dissolve any bases before adding them to solutions with ferrous iron; my experience is that adding solid potassium hydroxide to such solutions results in the formation of iron hydroxide around the pellet/flake which then oxidises to give ferric iron)

The resulting solution was a very dark yellow/green:

I left it in an amber vacuum desiccator with calcium chloride for about a couple of days; however after taking it out, I saw no change in the solution volume. Hence, I added ethanol as per the previous procedure to precipitate the complex. Unlike ferrous ammonium citrate, however, the potassium ferrous citrate complex merely formed a liquid bilayer:

Repeated mixing with a spatula and leaving it alone for a day did not result in any crystallisation; hence I decanted as much of the ethanol as possible. Interestingly, the ethanol is slightly orange, which suggests the possibility that ferric iron or a ferric complex dissolved into the ethanol.

(a few dark green droplets of the water/complex layer can be seen at the bottom)

I then boiled the complex down on a hotplate with added metallic iron as a reducing agent. At first, I got a dark, olive-green solid which then became a pale yellow-green upon further drying:

Comparing the tetrapotassium ferrous citrate (on right) to the previously prepared ferrous ammonium citrate (left):

Similar to the ferrous ammonium citrate, a dilute solution of the complex is a dark yellow-green:

Due to the colour difference vs freshly prepared potassium ferrous citrate, I'm not 100% certain as to the stochiometry or even if the compound is a citrate. It might be the case that the iron(II) ion was oxidised by air to iron(III), and was then reduced by citrate back to iron(II) with the citrate being oxidised to beta-ketoglutarate. I will note that the ferrous ammonium citrate and potassium ferrous citrate both resemble pictures of sodium ferrous citrate (see here and here). However, in my experience, industrial producers of such supplements tend to be good at guaranteeing the amount of the supplemented nutrient e.g. iron (and if you have a really good manufacturer, the amount of iron(II) or iron(III)), but not necessarily other factors such as whether the citrate is oxidised. Hence, I'm not sure if these pictures are a good comparison.

KoiosPhoebus - 17-3-2024 at 02:17

Recently, I found out that my ">90% potassium hydroxide" is actually closer to being 87-88% by acidimetry. Hence, I decided to repeat this synthesis with a slightly higher mass of potassium hydroxide (basically using 56.106 g/mol ÷ 0.88 as my "molar mass").

Interestingly, while the colour of the final solution remained very similar (a very dark yellow-green), the precipitate obtained by mixing with ethanol was instead a somewhat-pale yellow in colour:

Placing the new potassium ferrous citrate next to the previous batches (old on left, new on right):

This makes me wonder if the previous batch was stoichiometrically deficient in potassium, and mixing the solution with ethanol resulted in a product with at least one mono-protonated citrate. It would be an explanation for why the solution changed colour upon being mixed with ethanol.

Another possibility is that the increased alkalinity of the solution helped to deprotonate the hydroxyl group on citric acid. While the hydroxyl group should be difficult to ionise in aqueous solution (pKa > 11, depending on source), this publication describes a number of X-ray crystallography studies which demonstrate that citrate can coordinate metals through the hydroxyl group. A separate review describes the downward shift of pKas on nucleobases in the presence of metals which illustrates the possibility that the hydroxyl group is more readily ionisable in the presence of a metal which can compete with the hydrogen for binding to the oxygen.

A third possibility is that some sort of redox chemistry occurred at the higher pH. For example, maybe the iron(II) more readily oxidised to iron(III) and was reduced by the citrate in the presence of light. This would have resulted in the formation of 3-oxoglutarate which may have resulted in a different coloured complex. I should add here that adding the new potassium ferrous citrate to sodium carbonate did not result in the formation of a dark-green precipitate like with the old potassium ferrous citrate. Instead it formed a dark yellow solution which was somewhat closer to orange when compared to new potassium ferrous citrate solution. However, adding a few drops of hydrogen peroxide to a solution of the new potassium ferrous citrate did cause it to go orange-red, suggesting that the iron in the complex remains at Fe²⁺ (i.e. oxidation of the iron is probably not the cause of the colour difference).

Personally, I think the most likely possibility is the first. With the second, Fe²⁺ forms a much weaker complex with citrate than Fe³⁺ (K_Fe(II) = 10^4.8 versus K_Fe(III) = 10^11.27) and hence I doubt Fe²⁺ would be able to displace the hydrogen from the hydroxyl group to ionise it for complexation. While the third is also plausible, the colour of the new and old potassium ferrous citrate solutions appear very similar prior to the addition of ethanol, meaning that the oxidation state should not be significantly different between both samples (the only possibility is a redox reaction upon addition of ethanol). Additionally, I do add a small amount of metallic iron when boiling the solutions down and hence the complexed iron should not oxidise that readily.