Sciencemadness Discussion Board - Copper Borate? - Powered by XMB 1.9.11 (Debug Mode)

Copper Borate?

Arthur Dent - 29-5-2011 at 08:05

It seems that I have stumbled upon a interesting but mysterious compound... As I was contemplating what to do with my recently synthetized copper carbonate, I was looking at various acids and then I thought... hmm, boric acid and copper carbonate should yield Copper Borate, right?

So I started looking through Sciencemadness and found little if any info on this compound. A thorough search of the interweb thingie resulted in vague information at best... it appears that there is very little information on the properties of this mystery compound!

So far, here's what I found:

Copper Borate, also known as Cupric Borate or Tricopper Diborate: B2Cu3O6
Molecular Weight: 308.256 g/mol
Atomic structure:

It appears that there's another form of Copper Borate with the formula:
Cu(BO2

2 (CAS number 393290-85-2) (B2CuO4

molar mass: 149.1656 g

Here's a video (with slighly annoying music) on the synthesis of Copper Borate:
http://www.youtube.com/watch?v=ZOIeJQwkH-g

Many pyrotechnics websites mention that it is an interesting compound to make greenish blue stars. Some sites mention that the compound can be prepared by the reaction of Copper Carbonate with a warm solution of boric acid. Those same sites also mention that aside from the basic info above, the chemical is virtually undocumented! I have not found any info as to melting point, solubility, MSDS... (sigh)

I found patent Info on a copper borate-based pigment:
http://www.freepatentsonline.com/3100718.pdf

Here's a supposed photo of synthtized copper borate powder:
http://commons.wikimedia.org/wiki/File:Copper(II)_borate.JPG

So I turned to my books and was equally disappointed... Not even a passing mention in the big Brauer's inorganic Chem.

I found some interesting data from a 2006 article in the Russian Journal of Applied Chemistry at http://resources.metapress.com/pdf-preview.axd?code=f3hwj308... This is only a preview of the first page but it hints at a more detailed description of various forms of Copper borates:

amorphous:
2CuO•B2O3•H2O
3CuO•2B2O3•nH2O (where n could be 1,2,4,5)

crystalline:
CuB(OH)4Cl, a mineral known as Bandylite
http://webmineral.com/data/Bandylite.shtml

[Ca2Cu(OH)4B(OH)4]2, a mineral known as Henmilite
http://webmineral.com/data/Henmilite.shtml
which appears to be "ungodly rare" according to many sites.

It also describes some other exotic mineral and synthetic borates complexes whose formulaes are way too long to write up and not really pertinent to this thread. IVe put a copy of this article on my site at http://www.progmontreal.com/science/coppbora/ZincCopBora.jpg

So in conclusion, please describe your experiences with this compound. I'd be curious to see if 1) I can synthetize that compound and 2) actually know what i've specifically prepared.

Robert

Picture

LanthanumK - 29-5-2011 at 13:12

Here is a picture of copper(II) borate made by the reaction given on the video with slightly annoying music. I did not make this from the video; at this time, I just threw every metal chloride I could find into borax solution, filtered the precipitate (if there was any), dried it, and photographed it.

Edit: I just uploaded the file here that you linked to at Wikipedia, by the way.

[Edited on 29-5-2011 by LanthanumK]

The WiZard is In - 29-5-2011 at 13:56

Quote: Originally posted by LanthanumK

Here is a picture of copper(II) borate made by the reaction given on the video with slightly annoying music. I did not make this from the video; at this time, I just threw every metal chloride I could find into borax solution, filtered the precipitate (if there was any), dried it, and photographed it.

. la Condensed Chemical Dictionary sez -

Copper metaborate

Derivation: Interaction of copper sulphate and sodium borate.

You can get all the sodium borate you ever be needing/wanting
in your local supermarket.

djh
----
Who else remembers - Twenty Mule-team Borax, and
their TV commercials?

Trivia - Taking an item off the shelve and bring it to a clerk was
patented years ago. We now call it a supermarket the
patented/first supermarket [chain] is still in business - it be
be called ............?

I know because I am a long term subscriber to American Scientist
and a fan of Henry Petroski's column.

Shopping By Design
Supermarkets, like other inventions, didn't just happen; they were designed, developed—and patented
Henry Petroski
American Scientist 93[6] 491 & ff. November-December 2005

If there is overwhelming interest, as a subscriber I can obtain
a text copy. I do not save back issues of American Scientist so I cannot scan it.

[Edited on 29-5-2011 by The WiZard is In]

LanthanumK - 29-5-2011 at 15:20

Because I can always made copper(II) chloride but not copper sulfate, I used copper chloride. Either should work similarly.

Arthur Dent - 30-5-2011 at 03:17

Well, I tried to make a bit of copper borate, so in a suspended solution of copper carbonate (with just enough water to keep the slurry wet and mobile), I added a supersaturated solution of boric acid (warmed until all of it dissolved) and mixed the two together with heat and magnetic stirring on the hotplate.

The resulting hot solution was filtered and washed with hot water to remove the excess boric acid, and the resulting precipitate looks in every way similar to copper carbonate.

Maybe i'll re-do the experiment with carefully weighted quantities to determine the molar mass of what has been synthetized, then again, I don't have a very precise balance...

Robert

LanthanumK - 30-5-2011 at 03:30

Add excess boric acid, filter, dry, and throw the precipitate in HCl or H2SO4 to see how it dissolves. If it is unchanged, then no reaction occurred.

woelen - 30-5-2011 at 05:04

I have severe doubts on the feasibility of the reaction between copper carbonate and boric acid. Boric acid is a VERY weak acid and most likely no reaction occurs with the carbonate.

You might have better luck with freshly prepared and well rinsed copper hydroxide (prepared from a solution of NaOH and CuSO4, not from CuCl2). The Cu(OH)2 should not be dried, it just should be precipitated, rinsed and then used in further experiments.

Add some excess boric acid to a suspension of copper hydroxide in water and see if the color changes. Do not heat, heating will convert the hydroxide to black much more compact and less reactive oxide.

But even here I am not sure whether this really works or not. But you could give it a try.

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

Adding a solution of CuSO4 to a solution of borax does not yield a pure copper borate. A solution of borax is quite alkaline and at best you'll obtain some basic borate, which will contain a lot of hydroxide in the precipitate as well (same problem exists with making copper carbonate from aqueous solutions).

[Edited on 30-5-11 by woelen]

blogfast25 - 30-5-2011 at 05:35

Quote: Originally posted by woelen

I have severe doubts on the feasibility of the reaction between copper carbonate and boric acid. Boric acid is a VERY weak acid and most likely no reaction occurs with the carbonate.

You might have better luck with freshly prepared and well rinsed copper hydroxide (prepared from a solution of NaOH and CuSO4, not from CuCl2). The Cu(OH)2 should not be dried, it just should be precipitated, rinsed and then used in further experiments.

Add some excess boric acid to a suspension of copper hydroxide in water and see if the color changes. Do not heat, heating will convert the hydroxide to black much more compact and less reactive oxide.

But even here I am not sure whether this really works or not. But you could give it a try.

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

Adding a solution of CuSO4 to a solution of borax does not yield a pure copper borate. A solution of borax is quite alkaline and at best you'll obtain some basic borate, which will contain a lot of hydroxide in the precipitate as well (same problem exists with making copper carbonate from aqueous solutions).

[Edited on 30-5-11 by woelen]

My thoughts precisely, woelen.

Boric acid is an acid in name only, it is likely to be too weak to attack a water insoluble carbonate like CuCO3. Freshly precipitated Cu(OH)2.nH2O, maybe…

Likewise with the reaction between a neutral cupric salt and a borax solution: the high alkalinity of the borax solution is likely to precipitate either Cu(OH)2, a basic cupric salt or basic cupric borate. But it’s hard to tell which, without a solubility product (Ksp

for the target cupric borate.

Best chance to make a cupric borate is to fuse boric acid with CuO (or possibly with CuCO3, which decomposes at 200 - 300 C anyway) in at least Bunsen flame heat.

Arthur Dent - 30-5-2011 at 12:07

Yeah, i guess you're right, a brave attempt on my part, but going nowhere!

Plus I have to admit, I was just trying the cupric borate thing because I had a bit of leftover copper carbonate at the bottom of the beaker and just wanted to wash it off with some acid on hand. I'll just clean the whole mess and make like it never happened. Pretend you saw nothing! LOL

Robert

The WiZard is In - 30-5-2011 at 12:35

Quote: Originally posted by woelen

Adding a solution of CuSO4 to a solution of borax does not yield a pure copper borate. A solution of borax is quite alkaline and at best you'll obtain some basic borate, which will contain a lot of hydroxide in the precipitate as well (same problem exists with making copper carbonate from aqueous solutions).

Granted this is good enough and not necessarily CP.

http://tinyurl.com/4xjrran

----
From the Condensed Chemical Dictionary - 1919 via
Google.com/books

Copper Borate* (Cupric borate) CuBCu. Color and properties: Bluish-green, crystalline powder.

Soluble in water.
Derivation: By the interaction of copper hydroxide and boric acid.
Method of purification: Crystallization.
Grades: Technical.
Containers: Kegs; tins.
Uses: Oil pigment; painting on porcelain.
Fire hazard: None.
Railroad shipping regulations: None.

----------
25. Borate of copper. When borax is poured into a solution of
sulphate of copper, borate of copper is precipitated in the form of a
pale light-green jelly, which when dried is with great difficulty
soluble in water. It easily melts into a darkred vitreous
substance. According to Palm, by long trituration of filings of
copper and boracic acid in water, and then digesting the mixture, it
dissolves, and crystals may be obtained from it.

http://tinyurl.com/3j6788r

Hummm it sez most everywhere it is soluble not in water.

---
And here > X < I gave up looking.

Texium - 6-9-2014 at 16:54

Earlier I was looking for something interesting to do and decided to try to make copper borate using a solution of boric acid in 3% hydrogen peroxide on clean copper wire, to see if such a reaction is feasible. No blue color has appeared yet, but it hasn't been very long. Promisingly, it is bubbling very slowly, so there does seem to be a reaction happening. I'll report later if anything happens.

Update, 3.5 hours later: I just checked on it again, and now the solution is bubbling quite vigorously, but the solution is still clear and colorless. The copper still looks clean. I'm not sure exactly what is happening. I would have expected to see a blue color by now, either in solution or as a precipitate.

[Edited on 9-7-2014 by zts16]

deltaH - 6-9-2014 at 23:44

@zts

As for your experiment of copper wire in H2O2/boric acid:

The bubbling you are seeing is most likely the catalysed decomposition of hydrogen peroxide by the reaction:

2H2O2(aq) => O2(g) + 2H2O(l)

...and probably not much else

If you can collect the gas, you can test it. A glowing splinter lowered into a test tube filled with O2 should cause it to ignite.

****
To expand on the topic and as mentioned by others here, I do not think this preparation is very straight forward, but let me give my two cent suggestion on possibly preparing copper borates...

First let me say that borax is a quite basic salt that forms weakly alkaline solutions (~ pH 9) because boric acid is such a terribly weak acid. So, adding a solution of borax to say an acidic copper salt solution like CuSO4 or CuCl2, as others had originally done on this thread, probably simply precipitates copper hydroxide... hence the cyan coloured product? This can be tested as heating this product should cause it to turn black by forming cupric oxide through loss of water.

Possible preparation of copper borates by fusion:

One might have more success forming copper borates by fusing (melting) boric acid and precipitated copper hydroxide... maybe. At red heat, one might form a metaborate. One would expect such a material to be intensely coloured deep blue?

First the melt has to lose a lot of water, so expect much bubbling and formation of glassy viscous stuff, so be careful! Small amounts in a test tube over a flame with tongs would be my choice

Simply melting borax with copper hydroxides or oxides might form strongly coloured brittle borate glasses containing much copper. These could be interesting as pigments when ground up.

Hydrothermal preparation:

zts, if you're keen on experimenting with copper borates, another thing to try might be to dissolve copper hydroxide by adding excess ammonia solution (obtaining the deep blue copper ammine complex), then boiling this with boric acid in a fume cupboard or outside. Maybe this might give a copper borate too.

Bear in mind, my suggestions are purely hypothetical suggestions for experimentation, i.e. I do not know in advance that this would work!

****
As I mentioned, I would expect true copper borates to be deep blue insoluble compounds that do not turn black on heating.

[Edited on 7-9-2014 by deltaH]

Texium - 7-9-2014 at 07:24

Thanks deltaH, I'll try both of the methods that you mentioned at some point. Even if they don't work, they sound interesting.
And yeah, I had the thought that the peroxide was just decomposing in the back of my mind, but I kind of shut it out because I was really hoping that something would happen.

bismuthate - 7-9-2014 at 08:19

Wouldn,t the Cu(OH)2 decomose before reacting?

deltaH - 7-9-2014 at 08:42

Quote: Originally posted by bismuthate

Wouldn,t the Cu(OH)2 decomose before reacting?

Yes, both would probably dehydrate first at relatively low temperature, the hydroxide forming black cupric oxide and the boric acid forming metaboric acid, but hopefully, cupric oxide forming in a glassy melt of metaboric acid would be more amorphous than calcined alone and so [hopefully] more reactive. By the time you get to red heat, I am hoping it would react.

When I was a teenager, I used to flame fuse a lot of transition salts with borax including copper sulphate, I got a strongly coloured glassy bead which if I remember correctly was deep blue, but it was a very long time ago! I imagine that using boric acid should also work with obtaining some kind of borate powder, probably copper metaborate.

The colour changes should be: cyan => black =>[strong heating] =>>> deep blue

If anyone is interested, what I did back then was take a thin copper wire, make a tiny loop at the end (like the eye of a needle), scoop/dip that into a mixed grind of borax and copper sulphate and then place that in flame to get it red hot, when the powders melt, they form a bead on the tip of the wire and you can quickly 'glue' some more by dipping it in the powder while hot and glassy and repeat, thereby growing the bead to a small round coloured ball. Again, I can't be 100% sure about the colour, but think it was a dark deep blue when fully fused at high heat (initially it turned black).

[Edited on 7-9-2014 by deltaH]

Texium - 7-9-2014 at 08:48

I just checked my flask with the copper/boric acid/hydrogen peroxide, and there actually is a small amount of blue precipitate now, and it looks too dark to be copper hydroxide. It's still bubbling. Maybe something actually happened? If it did actually work though, it would be a horribly inefficient reaction, as the bubbling is most definitely from H₂O₂ decomp.

Quote: Originally posted by deltaH

Quote: Originally posted by bismuthate

Wouldn,t the Cu(OH)2 decomose before reacting?

So I guess that means that the wet method you described earier is out. Oh well, I'll still try the fusion method and see if I can get the results that you described.

deltaH - 7-9-2014 at 09:04

I wouldn't discard the wet method before trying, I wouldn't say it's out as having the [temporarily] stabilising ammine ligands may change things. The elevated pH of such a solution also changes the speciation of the borate ions in solution, possibly upping the concentration of borate ions significantly (as opposed to simply dissolved boric acid) and so this may help to form some kind of cupric borate. Anyhow, one can speculate till one is blue in the face (excuse the pun), but experiment will tell all, but bottom line, don't discard it until one tries it out experimentally.

****
Again, I want to re-iterate what woelen has already mentioned/alluded to above, the misleader with some of these methods is forming copper hydroxide only or mostly and this is probably what these cyan coloured precipitates are. If you heat it up and it turn black, it probably was just copper hydroxide or mixture thereof.

[Edited on 7-9-2014 by deltaH]

bbartlog - 8-9-2014 at 12:11

I prepared some copper borate (though it might better be called copper metaborate, since the copper:boron ratio is 1:2). Wet chemistry works fine, no need for high temperature fusion of solids.

http://www.sciencemadness.org/talk/viewthread.php?tid=16586&...

deltaH - 10-9-2014 at 22:25

Interesting bbartlog, but I would still try a fusion and perhaps the ammonia complexation method because there may be many types of 'copper borates' to prepare.

Personally, I do not believe you prepared a metaborate, for example, not because of the stoichiometry, but because metaboric acid forms at higher temperature by dehydration. From Wikipedia the rough profile of heating and dehydrating boric acid is:

Boric acid => [T > ~170C] => metaboric acid => [T > ~300C] => higher polyborates and on continued heating, eventually boron trioxide

Also, the tetrahydroxymonoborate anion, B(OH)₄^-, could also hypothetically give a copper borate with a ratio of 1:2 copper to boron, although I would have expected this particular borate ion to dominate only at high pH and so am surprised if this formed by raising the pH with bicarbonate only.

This is why I proposed complexing/stabilising the copper with ammonia first, as it allows one to access much higher pH's where the borate speciation would likely be more densely charged.

bbartlog - 11-9-2014 at 06:33

I expect you are right about it not being metaborate; that should really just be take as a comment on the apparent stoichiometry.

The B(OH)₄- ion doesn't need to dominate in order to be implicated, it just needs to form something really insoluble with copper, but the weight/stoichiometry for Cu++(B(OH)₄-)₂ doesn't work out unless some arbitrary weight of H₂O is removed from the result.

It occurs to me though that while the ratio of boron to copper can't be greater than 2 in my precipitate (because no copper remained in solution), it could easily be less than 2, for example if there is some stable copper hydroxyborate along the lines of Cu-OH-B(OH)₃. I didn't do any checking for that possibility. If I repeat the experiment with a smaller (50%) quantity of boric acid and weigh the result carefully I should be able to see whether that is a plausible explanation.

Shivachemist - 5-2-2015 at 00:33

Well, metaborates are mostly more alkaline than tetraborates. Usually, carbonates and bicarbonates are used to prepare soluble tetraborates and Borax for precipitating insoluble tetraborates. I can see many members had discussed about the reaction between Boric acid and carbonates. A warm concentrated solution of Boric acid will definitely attack carbonates and bicarbonates and produces corresponding tetraborate with the evolution of carbon dioxide but rather slowly. You can even check it by shaking boric acid with cold sodium carbonate solution. I had tried it many times and it worked exceedingly well. I have even made videos about Ammonium & Sodium tetraborate. But, there are some difficulties in making Ammonium tetraborate, I explained that in my video.

Here is the link of my recent video(Preparation & Properties of Ammonium tetraborate)

https://www.youtube.com/watch?v=iKd1e_qJHuY

As far as metaborates are concerned, they can be obtained by treating the hydroxides with Boric acid. But, this only works for weak bases. For strong bases, you need to add polyhydroxyl compounds such as Mannitol, Sorbital, Glycerol etc.. to increase the acid strength and then you can titrate with strong bases.

blogfast25 - 5-2-2015 at 08:07

It's a nice video Shiva but Id still like to see more evidence that your product is indeed ammonium borate, which borate and which hydrate, On those points you are rather vague.

As regards:

Quote:

For strong bases, you need to add polyhydroxyl compounds such as Mannitol, Sorbital, Glycerol etc.. to increase the acid strength and then you can titrate with strong bases.

... what evidence/references do you have?

[Edited on 5-2-2015 by blogfast25]

Shivachemist - 5-2-2015 at 10:30

Well, thank you. You can even see this quoted point in wiki too (see under reactions in wiki). Btw, if you wanna learn something new or do something new in the subject you like, you should develop from the basics and add your own part too to show something new. We can't depend on the references all the time. A compound is not in the wiki does not mean, it does not exist at all. I can prove the nature of a chemical by doing experiments, but wiki not gonna accept research based points too because they had told me so. I do not wanna copy other people points all the time, I just wanna learn something from what others have written and do something new. For this reason, I quit writing in wiki a long time back.

If someone has done something new, whether it is correct or not, it does not matter, trying something new is important, that's how we can grow and improve our knowledge. If you have doubts in it, you can ask him to clarify it by asking questions, if he has references for it, it is fine, but we can't have references all the time. You can learn something from a book and try something differently. I have seen many are saying, this is in wiki, that is not in wiki and all, well, wiki is not the only platform to learn. It is one of the big platforms to learn things, I'm not denying it because I have learned so much from wiki too. Sometimes you can find a way to prepare a compound that has not been mentioned anywhere online or even in books.

Well, the one I prepared passed the following tests:

(i) Borate flame test (pic attached)
(ii) Precipitates Boric acid with cold solution of ammonium borate on adding dilute sulfuric acid (pic attached)
(iii) Liberates ammonia when heated with alkali.

Also, I wanna mention the difference between borates and metaborates.

For example, Sodium metaborate is more alkaline than Borax. Its aqueous solution behaves like a mixture of borax and hydrogen peroxide. For this reason, it is usually used as a mild oxidizing and bleaching agent. On heating, it will liberate oxygen.

[Edited on 5-2-2015 by Shivachemist]

[Edited on 5-2-2015 by Shivachemist]

blogfast25 - 5-2-2015 at 11:26

Quote: Originally posted by Shivachemist

Well, thank you. You can even see this quoted point in wiki too (see under reactions in wiki). Btw, if you wanna learn something new or do something new in the subject you like, you should develop from the basics and add your own part too to show something new. We can't depend on the references all the time.

Thanks for the lecture (yawn!)

Relying on what others have done before and reported in peer reviewed sources is part (a BIG part!) of the Scientific Method.

Asking for references of prior art/existing evidence is too.

Without knowing prior art, what you currently know about chemistry would fit on the back of a small stamp. Don't knock it.

Shivachemist - 5-2-2015 at 11:57

You did not get my point I think, fine, no worries. Sorry if I've not made it clear. Thanks for reading it though. But, I'm glad I got the opportunity to interact with you all :-) Thanks once again.

blogfast25 - 5-2-2015 at 13:17

Oh, I got your point alright, don't worry.

[Edited on 6-2-2015 by blogfast25]

DraconicAcid - 5-2-2015 at 17:59

Quote: Originally posted by blogfast25

Quote:

For strong bases, you need to add polyhydroxyl compounds such as Mannitol, Sorbital, Glycerol etc.. to increase the acid strength and then you can titrate with strong bases.

... what evidence/references do you have?

Cotton & Wilkinson, Advanced Inorganic Chemistry, 4th Ed, p 298: In the 1:1 complexes the acidity of the OH groups exceeds that of B(OH)3, so that if glycerol is added to a boric acid solution, this can be titrated using aqueous NaOH.

blogfast25 - 5-2-2015 at 19:18

Quote: Originally posted by DraconicAcid

Thanks, DA. New to me, that bit.

Simple question, simple answer, no lecture. Just like I like it.

[Edited on 6-2-2015 by blogfast25]

DraconicAcid - 6-2-2015 at 09:07

Quote: Originally posted by blogfast25

Thanks, DA. New to me, that bit.
Simple question, simple answer, no lecture. Just like I like it.

Aw, but I had almost finished typing up the lecture I was going to attach to it......

deltaH - 6-2-2015 at 09:29

Quote: Originally posted by Shivachemist

For example, Sodium metaborate is more alkaline than Borax. Its aqueous solution behaves like a mixture of borax and hydrogen peroxide. For this reason, it is usually used as a mild oxidizing and bleaching agent. On heating, it will liberate oxygen.
[Edited on 5-2-2015 by Shivachemist]

Are you not confusing sodium metaborate with sodium perborate?

Shivachemist - 6-2-2015 at 09:34

Well, it was formerly considered to be a perborate, NaBO3. 4 H2O ------> NaBO2. H2O2. 3 H2O

deltaH - 6-2-2015 at 09:52

Quote: Originally posted by Shivachemist

Well, it was formerly considered to be a perborate, NaBO3. 4 H2O ------> NaBO2. H2O2. 3 H2O

Do you have a reference to suggest that it's a perhydrate and the peroxide is not formally bonded to the boron?

I know that common 'sodium percarbonate' is incorrectly named and is, in fact, sodium carbonate perhydrate.

Shivachemist - 6-2-2015 at 10:13

Well, yes, page 2.236 in the book "Textbook of Inorganic chemistry" written by PL Soni and Mohan Katyal.

This is the third time me giving references to my points. I'm really tired of this. This is ridiculous. This is more like a student is asking references for whatever the teacher is saying just because he does not know it. If you believe me or not, it does not matter to me. I'm just giving my suggestions and saying points that I have read. Believing it or not is up to you. Asking questions to clarify the doubts is different and asking references for everything is different. I had worked as a moderator in a forum too for a long time and written thousands of posts and I had answered many members questions too, at least from what I have learned heretofore. But, none has asked me references for whatever I'm saying. May be this is your practice, I do not know. If it so, then I'm not up for this.

Thank you.

blogfast25 - 6-2-2015 at 10:14

Quote: Originally posted by DraconicAcid

Aw, but I had almost finished typing up the lecture I was going to attach to it......

Ah yes Sir, but do you have a license to write lectures on this forum?

http://youtu.be/WnlIWpZSPXU

[Edited on 6-2-2015 by blogfast25]

Shivachemist - 6-2-2015 at 10:16

oh well then, I will stop lecturing and you continue it.

blogfast25 - 6-2-2015 at 10:21

Quote: Originally posted by Shivachemist

This is the third time me giving references to my points. I'm really tired of this. This is ridiculous. This is more like a student is asking references for whatever the teacher is saying just because he does not know it. If you believe me or not, it does not matter to me.

You really don't get it do you? Mostly it has nothing to do with 'believing you/not believing you' and only a matter of establishing where you obtained a particular bit of wisdom from.

Having said that, give me one good reason why I should blindly believe you (or vice versa, for that matter)... because you know everything and your knowledge base is beyond scrutiny?

[Edited on 6-2-2015 by blogfast25]

Shivachemist - 6-2-2015 at 10:30

I never think myself brilliant and all. I have a very ordinary knowledge in chemistry. But again, I'm just saying what I know. I can ask the same question back. Btw, did you ask your teacher references for whatever he/she is saying when you were studying? Just think. For saying this, do not tease me by calling sir and all, me just saying for example. If you do not want me to post my opinions, I will stop writing. I have no issues with it.

Thank you.

deltaH - 6-2-2015 at 10:31

Quote: Originally posted by Shivachemist

Well, yes, page 2.236 in the book "Textbook of Inorganic chemistry" written by PL Soni and Mohan Katyal.

This is the third time me giving references to my points. I'm really tired of this. This is ridiculous. This is more like a student is asking references for whatever the teacher is saying just because he does not know it. If you believe me or not, it does not matter to me. I'm just giving my suggestions and saying points that I have read. Believing it or not is up to you. Asking questions to clarify the doubts is different and asking references for everything is different. I had worked as a moderator in a forum too for a long time and written thousands of posts and I had answered many members questions too, at least from what I have learned heretofore. But, none has asked me references for whatever I'm saying. May be this is your practice, I do not know. If it so, then I'm not up for this.

Thank you.

The only reason is because I was giving you the benefit of the doubt. The wiki article states that it is the borate anion that carries the peroxide, so it should formally be named perborate. Your claim that it is, in fact, a perhydrate goes contrary to the wiki article which claims the peroxide is bonded to the boron, so simply wanted to verify which is correct.

blogfast25 - 6-2-2015 at 10:43

Quote: Originally posted by Shivachemist

Btw, did you ask your teacher references for whatever he/she is saying when you were studying? Just think.

If you do not want me to post my opinions, I will stop writing. I have no issues with it.

Your first point is a silly analogy: students aren't scientists yet. They need to pass exams based on the existing curriculum. In a good education system scepticism can be encouraged though.

Your second point is childish: no one wants you to stop posting here.

Take a step back, a deep breath and come back when you can see clearer.

[Edited on 6-2-2015 by blogfast25]

Chemosynthesis - 6-2-2015 at 11:24

I know my opinion is uninvited, but I have enjoyed your posts thus far, shiva chemist, and would hope you wouldn't stop posting due to references. I was unaware of both of the previous claims you made in the thread, and learned them from your posts and the subsequent references. It can be mildly inconvenient to cite things, but I try to have citations available for any claims I make here, or in general unless otherwise qualified. Personally, I find it very instructive as it can provide context on when literature is incorrect, or how long something has been known.

As to teachers, etc. it is commonplace for colleagues I know, myself included, to ask each other for citations on claims when conversing, particularly regarding research, and many of these being asked are considered world-renowned experts in their fields, some with multiple doctorates. Graduate students and post docs are encouraged to question everyone as well, and this can be a very good source of constructive criticism for grant review, since sometimes they are less indoctrinated in the predominant groupthink.

Bert - 6-2-2015 at 12:02

Please everyone, just relax. We come from different cultures.

It's standard to give reference, if you've already done so and someone asks, just say "see further up thread", or drop a link to earlier post.

blogfast25 - 6-2-2015 at 13:35

Quote: Originally posted by Bert

We come from different cultures.

With all due respect but I don't see what that has to do with anything. Shivachemist (or anyone else here) is hardly representative of a culture. An individual can't be that.

[Edited on 6-2-2015 by blogfast25]

quantumcorespacealchemyst - 10-2-2015 at 19:59

thanks for the video Shivachemist and the flame test + other results. i am interested in this copper borate/metaborate topic.

i saw recently (http://silver.atomistry.com/silver_borate.html that reportedly) "A solution of borax reacts with one of Silver nitrate to precipitate the white metaborate, AgBO2. It is also produced by dissolving Silver monoxide in boric acid, an equilibrium being attained. Conversely, water causes partial hydrolysis of silver borate to Silver monoxide and boric acid.". it doesn't seem to have a source.

i find it interesting that Copper nitrate makes hydrates, and decomposes at 180°C whereas Silver nitrate can be heated and not decompose and doesn't seem to form hydrates (does it?), but it's borate supposedly dissolves/decomposes in water, where copper borate doesn't (i believe that is the consensus above).

blogfast25 - 11-2-2015 at 10:22

Quote: Originally posted by quantumcorespacealchemyst

thanks for the video Shivachemist and the flame test + other results. i am interested in this copper borate/metaborate topic.

i saw recently (http://silver.atomistry.com/silver_borate.html that reportedly) "A solution of borax reacts with one of Silver nitrate to precipitate the white metaborate, AgBO2. It is also produced by dissolving Silver monoxide in boric acid, an equilibrium being attained. Conversely, water causes partial hydrolysis of silver borate to Silver monoxide and boric acid.". it doesn't seem to have a source.

i find it interesting that Copper nitrate makes hydrates, and decomposes at 180°C whereas Silver nitrate can be heated and not decompose and doesn't seem to form hydrates (does it?), but it's borate supposedly dissolves/decomposes in water, where copper borate doesn't (i believe that is the consensus above).

Atomistry rarely provides verifiable references.

Abromination - 11-7-2018 at 13:12

I know its been a while, but I managed to get borax and copper sulfate to react with eachother and not form something that looks like copper hydroxide. Using an excess of copper sulfate, I added borax powder to a concentrated-ish solution of copper (ii) sulfate. I filtered the jelly like precipitate and dried it. I rinsed off excess copper sulfate and was left with dark blue/green chuncks.

Lion850 - 2-9-2020 at 00:26

I also had a go at copper borate; specifically hoping for copper metaborate Cu(BO2)2.

First attempt:
- 15g sodium metaborate NaBO2 dissolved in 100ml water, heated until dissolved in a clear solution.
- 20g copper sulphate pentahydrate dissolved in 100ml water, heated to dissolve all.
- Slowly add the copper sulphate solution to the sodium metaborate solution (both hot). Initially there was just a slow color change but suddenly there was a lot of gas evolution and the beaker bubbled over. The gas had no smell. This was unexpected, still don't know why it happened.
- Solution turned baby blue.
- Stir for 60 minutes and vacuum filter. Clear filtrate, baby blue remainder. Photo below. Rinse once in funnel with hot water.

- Dry on steam bath until weight stable. Color became slightly darker. Photo below Final recovery 10g.

The product dissolved quite easily in concentrated HCl giving as yellow-ish solution, with tiny bit of unreacted material. I also heated some product in a test tube on a bunsen burner for 10 minutes; it turns more green. Photo below. The green color remains when the crystals are wetted.

I tried to do a flame test; there is a hint of green - much less green than when I do a test with boric acid.

As mentioned a few times in this thread, what I got was probablky a mix of a copper borate and copper hydroxide?? I had another go while trying to reduce the boric solution prior to adding the copper sulphate.

Second attampt:
- 20g copper sulphate pentahydrate in 120ml water, room temperature. Stir till all dissolved. I purposely used the same amount of CuSO4.5H2O.
- 15g sodium metaborate in 120ml water, stir to dissolve at room temperature. pH very high, 11 or more.
- Add boric acid H3BO3 to the NaBO2 solution and keep checking pH, stop after 5g added because some boric acid did not dissolve. pH around 9 as far as I can tell from the strips.
- Mix slowly - very slowly - with both solutions at room temperature.
- Some gas evolution, for a few seconds, but much less than with the first attempt.
- Baby blue solution, similar color as first attempt.
- Stir (no heating) for 24 hour. Color remains the same, see photo below. No more unreacted boric acid to be seen.

- Vacuum filter. Wash once in funnel with large volume of room temp water.
- Dry on steam bath until weight stable. See before and after photos below; dried product is much darker blue than the first attempt. Final recovery 12g.

Attempt at a flame test shows very little green.
When heated on the bunsen burner, this product also turns green, photo below.

After this, I wanted to make copper hydroxide and dry it using the same steam bath to see how the final product compares to the above two. This went as follows:

- A solution containing 20g copper sulphate pentahydrate was added to a solution of 8g sodium hydroxide when both solutions were warm - a blue suspension formed but turned black within seconds. Looked like the copper hydroxide decomposed to copper oxide, and this was discarded.

After more reading up, the reaction was tried again but with much more dilute solutions:
- 20g CuSO4 dissolved in 400ml water
- 8g NaOH dissolved in 100ml water.
- Mix the room temperature solutions slowly, initially a light blue suspension which became a lovely dark blue. See photo. Solution was quite thick and hard for the stir bar to get moving.

- Vacuum filter and rinse once in funnel. Clear filtrate, dark blue remainder. See photo. The remainder retained a lot more water than the "borates" and took much longer to dry.

- As it dried the color changed became darker and eventually alomost black, see photo. I suspect this became a mix of hydroxide and oxide.

- Final recovery 9g.
- The product dissolves readily in contentrated HCl, but a bit slower than the "borates".

Thus the 'borates' are different from the hydroxide; they are much more stable in terms of keeping their color at the steam bath temperature, and seem to retain less water i.e. they dry easier. But what I really got I dont know.

I'm tempted to have a go at stirring copper powder with boric acid and see what happens, but these kind of reactions ties up my (only) stirring hotplate for days. So it will have to wait as something else is currently on the plate.

Photo for comparison:

Bedlasky - 2-9-2020 at 03:58

I think that blue precipitate is some polyborate (like di-, tri-, tetraborate etc.) and the green stuff if actually metaborate. You can see some other black material in your green solid - this is probably CuO.

unionised - 2-9-2020 at 05:17

one of the earlier references in this thread

https://books.google.co.uk/books?id=MGQ9AAAAYAAJ&pg=PA66...
says that copper borate melts easily .
Copper (hydr)oxide won't melt easily.

Are you able to check if your product forms a dark red glassy material when melted + cooled?

[Edited on 2-9-20 by unionised]

Lion850 - 2-9-2020 at 17:36

Gents thanks for the replies. I was aware of that report of copper borate melting 'easily' but when I heated my 2 products in a strong bunsen burner flame they only turned green.

Clear info on what happens to copper borates when heated seems not that easy to find online. The attached document describes synthesis of copper borates and has a section on their thermal decomposition:

"The temperature of the thermal decomposition process of copper hydroxyl nitrate was reported to be 246 °C [9]
which was very close to those of the precipitates obtained in the present study. The mass loss onset temperatures of the all samples were between 237.7 and 248.7 °C as shown in Table S1. The remaining mass after thermal decomposition would be 66% considering the stoichiometry of the reaction in Eq. 2. The remained mass from the samples at 600 °C were in the range of 63.4% and 70.9% as indicated in Table S1."

It does not mention melting, rather decomposition as far as I can make out.

Attachment: Prep of copper borates as lubricant additives.pdf (537kB)
This file has been downloaded 267 times

Attempt at prep. of copper borate

old school - 28-9-2020 at 04:10

Forgive me for not posting exact weighed measurements, but I added about 2 tablespoons of copper sulphate "feed grade" crystals in about 250 ml of distilled H2O in one flask, and then ~ 2 tablespoons of laundry borax in ~250 ml of distilled H2O in another flask, warmed both both flasks for ~1 minute in microwave, stirred well. Then dumped the contents of the copper sulphate flask into the borax flask, stirred by hand another minute or two.

Next, poured it all into a teflon beaker and placed on a magnetic stirrer hot plate stirring vigorously, ran the temp. up to 62C, kept it for 2 hours. Used a simple coffee filter to filter out the precipitate, washed the cuprous borate out several times with distilled H2O over an hour, it drained very, very slowly - - a vaccum filtration would have been better. Finally, placed the turquoise filter cake on a teflon mat in the oven drying it at ~93C for 2 hours. Disposed of the Na2So4 filtrate.

If there was any CuOH2 in the final product it would have been revealed by the rather classic lab procedure of heating a small portion of my final product in distilled H2O at ~97C, just below boiling for an hour which converts any CuOH2 present to black CuO, with a few boiling stones for good measure. I did that.

Result: No black color was observed at all.

From the patent literature:
For example, 763 parts of water solution containing 10 percent decahydrated borax Na₂B₄O₇ · 10 H₂O were added, under stirring, to 125 parts of water solution containing 40 percent pentahydrated copper sulphate CuSO₄· 5H₂O at 50 oC. The solution was kept at 50 oC for 1 hour, and then kept at 20 - 30 oC for one night, consequently resulting in a precipitation. The precipitation was filtered and washed with water repeatedly until unreacted borax was completely removed, and dried at 40 ° C. Consequently, 26.81 grams of copper tetraborate,dibasic were obtained.
3. COPPER TETRABORATE TRIBASIC--3Cu(OH)₂·CuB₄O₇--
[0027]
Copper tetraborate,tribasic is prepared by mixing a solution containing a high concentration of copper sulphate and a solution containing a high concentration of borax at a remarkably high temperature, e.g., 60 oC with adding a greater amount of water with maintaining a ratio of one mole copper sulphate to one mole borax. Specifically, after copper tetraborate is produced as shown in FORMULA (1), water washing is executed repeatedly to produce copper tetraborate,tribasic as shown in FORMULA (4).
Figure imgb0003
[0028]
For example, 12.5 parts of water solution containing 40 percent pentahydrated copper sulphate CuSO4· 5H2O - - - - - and - - - -
75 parts of water solution containing 10.2 percent decahydrated borax Na₂B₄O₇ · 10 H2O were mixed with maintaining the ratio of 1 to 6, i.e., the ratio of one mole to one mole, at 60 ° C.
100 parts of water were added into the solution at 60 °C, and was kept at 60 °C for 1 hour, and then kept at 20 - 30°C for one night, consequently resulting in a precipitation. The precipitate was filtered and washed with water repeatedly until unreacted borax was completely removed, and dried at 60° for 48 hours. Consequently, 2.43 grams of copper tetraborate, tribasic were obtained.

https://patents.google.com/patent/EP0450568A2/en

Lion850 - 29-9-2020 at 21:07

Hi old school can you post a photo of your final product? I am interested to see how the color compares to what I go as per my post a bit further up.

AJKOER - 30-9-2020 at 07:49

I just have an idea which I am now investigating. It is based on magnesium borate as a path to say tetraamine copper borate, which on mild heating may successfully release NH3, leaving copper borate.

Paths to magnesium borate (see https://www.sciencedirect.com/science/article/abs/pii/S00223...), seemingly best: borax (Na2B4O7) + Epsom Salt (MgSO4) + freeze to produce a Na2SO4 hydrate, leaving MgB4O7.

Idea is based generally on the action of NH4X on Cu metal in the presence of O2/H2O2 and a tiny bit of NaCl is an electrochemical path to [Cu(NH3)x(H2O)y]Xz where x + y = 6 and z = 1 or 2 (see comments and references here http://www.sciencemadness.org/talk/viewthread.php?tid=90354#... ).

Just employing Cu/O2/NH3(aq) expected product is [Cu(NH3)x(H2O)y](OH)2 (or the corresponding cuprous).

Add the Magnesium borate to precipitate Mg(OH)2 leaving a corresponding copper aqua amine borate (which may have a commercial application as a fungicide, see 'LABORATORY EVALUATION OF BORATE:AMINE:COPPER DERIVATIVES IN WOOD FOR FUNGAL DECAY PROTECTION' at https://www.fpl.fs.fed.us/documnts/pdf2011/fpl_2011_chen001....). Mild heating leaves Copper borate.

Advantages: All inexpensive and available reagents.

Before I proceed, any suggestions?

[Edited on 30-9-2020 by AJKOER]