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[*] posted on 29-5-2011 at 08:05
Copper Borate?


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: B<sub>2</sub>Cu<sub>3</sub>O<sub>6</sub>
Molecular Weight: 308.256 g/mol
Atomic structure:


It appears that there's another form of Copper Borate with the formula:
Cu(BO<sub>2</sub>;)<sub>2</sub> (CAS number 393290-85-2) (B<sub>2</sub>CuO<sub>4</sub>;) 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•B<sub>2</sub>O<sub>3</sub>•H<sub>2</sub>O
3CuO•2B<sub>2</sub>O<sub>3</sub>•nH<sub>2</sub>O (where n could be 1,2,4,5)

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

[Ca<sub>2</sub>Cu(OH)<sub>4</sub>B(OH)<sub>4</sub>]<sub>2</sub>, 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




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[*] posted on 29-5-2011 at 13:12
Picture


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.

Copper(II) borate.JPG - 69kB

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

[Edited on 29-5-2011 by LanthanumK]
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[*] posted on 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]
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[*] posted on 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.
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Arthur Dent
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[*] posted on 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




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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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 (K<sub>sp</sub>;) 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.
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[*] posted on 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




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[*] posted on 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.

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[*] posted on 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]




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[*] posted on 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]




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[*] posted on 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. :(




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[*] posted on 7-9-2014 at 08:19


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



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[*] posted on 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]




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[*] posted on 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?


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.
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.



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[*] posted on 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]




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[*] posted on 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&...







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[*] posted on 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.







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[*] posted on 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.




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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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.

IMG_0956.JPG - 1.1MB IMG_0957.JPG - 2MB

[Edited on 5-2-2015 by Shivachemist]

[Edited on 5-2-2015 by Shivachemist]
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[*] posted on 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.




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[*] posted on 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.
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