Sciencemadness Discussion Board

Copper sulfate using chloric acid

Amos - 8-10-2014 at 13:26

The traditional way most people seem to make copper(II) sulfate for copper chemistry uses either sulfuric acid and hydrogen peroxide, or electrolysis of sulfuric acid with copper electrodes. Since I can only obtain 3% H2O2 and don't have a sufficient apparatus for electrolysis, and because I also don't possess any nitric acid, an idea dawned on me. Chloric acid is highly oxidizing, so I decided to add about ~3 grams of potassium chlorate to 100ml of 10% sulfuric acid solution, tossed in some copper, and waited. After about 15 minutes, nothing happened as expected. But when I put the beaker on a hotplate just below the boiling temperature of water, I had a lovely blue solution of copper sulfate within 5 minutes! There don't appear to be any chlorinated gases coming off due to chloric acid decomposition; I suspect the dilute nature of the acid safeguards that. So now I have a way of making copper(II) sulfate in situ without electrolysis, the waiting game of hydrogen peroxide, or any nitrogen oxide fumes in my lab.

gdflp - 8-10-2014 at 13:32

Nice job! One problem with this method however is the introduction of potassium ion contamination, though this could be removed by precipitating with sodium carbonate filtering, then redissolving using more sulfuric acid. Also, wouldn't it be cheaper though to by root killer(though I have no problems with making it for the fun of making it)? I'm surprised you didn't get any chlorine or chlorine dioxide. I'm also guessing that using sodium or potassium nitrate could be a cheaper alternative to the chlorate, unless it is made cheaply by electrolysis;)

[Edited on 10-8-2014 by gdflp]

Little_Ghost_again - 8-10-2014 at 13:34

Quote: Originally posted by gdflp  
Nice job! One problem with this method however is the introduction of potassium ion contamination, though this could be removed by precipitating with sodium carbonate filtering, then redissolving using more sulfuric acid. Also, wouldn't it be cheaper though to by root killer(though I have no problems with making it for the fun of making it)? I'm surprised you didn't get any chlorine or chlorine dioxide. I'm also guessing that using sodium or potassium nitrate could be a cheaper alternative to the chlorate, unless it is made cheaply by electrolysis;)

[Edited on 10-8-2014 by gdflp]


Root killer here in Scotland just contains weedol type weed killer

Texium - 8-10-2014 at 13:39

I've found that using 3% hydrogen peroxide with acids still works fairly well for making copper compounds, it just makes the solution really dilute since you have to use so much of it. I've been looking for more concentrated stuff, but I haven't found any yet.

Amos - 8-10-2014 at 13:42

Quote: Originally posted by gdflp  
Nice job! One problem with this method however is the introduction of potassium ion contamination, though this could be removed by precipitating with sodium carbonate filtering, then redissolving using more sulfuric acid. Also, wouldn't it be cheaper though to by root killer(though I have no problems with making it for the fun of making it)? I'm surprised you didn't get any chlorine or chlorine dioxide. I'm also guessing that using sodium or potassium nitrate could be a cheaper alternative to the chlorate, unless it is made cheaply by electrolysis;)

[Edited on 10-8-2014 by gdflp]


I have bought root killer, but I've found it's actually cheaper where I am to use online bought sulfuric acid, considering I have plenty of copper lying around from recycled electronics. Using a nitrate produces nitric acid, so there you get nasty nitrogen dioxide coming off and your acid only lasts so long. I can't say for certain, but I think that the chloric acid just gets recycled, forming copper(II) chlorate and immediately being turned back into chloric acid by the sulfuric acid in solution.

gdflp - 8-10-2014 at 14:02

Something has to be reduced. The reaction of chloric acid and copper will reduce the chlorine to low oxidation states, possibly down to chloride(I'm not sure if hypochlorite is capable of oxidizing copper. I'm guessing your getting some chlorine containing gasses coming off the solution, especially since your heating it. I was under the impression that chloric acid is extremely prone to decomp on heating.

Amos - 8-10-2014 at 14:11

Quote: Originally posted by gdflp  
Something has to be reduced. The reaction of chloric acid and copper will reduce the chlorine to low oxidation states, possibly down to chloride(I'm not sure if hypochlorite is capable of oxidizing copper. I'm guessing your getting some chlorine containing gasses coming off the solution, especially since your heating it. I was under the impression that chloric acid is extremely prone to decomp on heating.


Well like I said, It was a dilute solution, and I never smelled anything, even once the solution was half of the original volume. I imagine any chloride ions would have presented themselves as green coloration of the solution, especially when hot, but it was only ever deep blue. Maybe it could've evolved hydrogen gas? My knowledge of the finer aspects of reaction chemistry aren't too great.

[Edited on 10-8-2014 by No Tears Only Dreams Now]

gdflp - 8-10-2014 at 14:19

I guess the sulfuric acid could be reduced, though then you would smell sulfur dioxide, and that also doesn't make sense because I believe chloric acid is a much better oxidizer than sulfuric acid.

blargish - 8-10-2014 at 14:41

Quote: Originally posted by gdflp  
I guess the sulfuric acid could be reduced, though then you would smell sulfur dioxide, and that also doesn't make sense because I believe chloric acid is a much better oxidizer than sulfuric acid.


Yea, I believe that sulfuric acid only showcases oxidizing properties in hot, concentrated solutions.

I am pretty sure the chloric acid is being reduced, it's just a question of what it's being reduced to. Chloride is probably the best bet, with its color being masked by the copper sulfate.

Amos - 8-10-2014 at 14:55

I'm getting some weird and very exciting results now, guys. I took my solution of copper sulfate/whatever else and poured it into a gross excess of sodium bicarbonate, precipitating copper(II) carbonate as usual. After I filtered the slurry, though, I still had a somewhat cloudy blue solution left. Figuring I must've not used enough of the baking soda, I added more. No change occurred, the baking soda just dissolved. I decided "let's try something stronger!", and poured about 5 grams of sodium hydroxide prills in, and to my complete and utter surprise, it became DARKER blue, and completely cleared! It's almost as deep as the tetraammine copper(II) complex now!

blargish - 8-10-2014 at 15:02

Quote: Originally posted by No Tears Only Dreams Now  
I'm getting some weird and very exciting results now, guys. I took my solution of copper sulfate/whatever else and poured it into a gross excess of sodium bicarbonate, precipitating copper(II) carbonate as usual. After I filtered the slurry, though, I still had a somewhat cloudy blue solution left. Figuring I must've not used enough of the baking soda, I added more. No change occurred, the baking soda just dissolved. I decided "let's try something stronger!", and poured about 5 grams of sodium hydroxide prills in, and to my complete and utter surprise, it became DARKER blue, and completely cleared! It's almost as deep as the tetraammine copper(II) complex now!


Would the dark blue have something to do with the formation of Cu(OH)2?

Amos - 8-10-2014 at 15:20

Quote: Originally posted by blargish  
Quote: Originally posted by No Tears Only Dreams Now  
I'm getting some weird and very exciting results now, guys. I took my solution of copper sulfate/whatever else and poured it into a gross excess of sodium bicarbonate, precipitating copper(II) carbonate as usual. After I filtered the slurry, though, I still had a somewhat cloudy blue solution left. Figuring I must've not used enough of the baking soda, I added more. No change occurred, the baking soda just dissolved. I decided "let's try something stronger!", and poured about 5 grams of sodium hydroxide prills in, and to my complete and utter surprise, it became DARKER blue, and completely cleared! It's almost as deep as the tetraammine copper(II) complex now!


Would the dark blue have something to do with the formation of Cu(OH)2?


No no no, I've spent a lot of time doing copper chemistry, copper(II) hydroxide is insoluble, it makes for cloudy, gross colored slurries. Also, this solution turns violet(blue+magenta) when phenol red indicator is added, indicated that it's strongly basic. Here's the solution on its own:

IMG_0926.JPG - 3.6MB

Brain&Force - 8-10-2014 at 15:26

It probably has something to do with the presence of chlorate. It reminds me of this:

http://woelen.homescience.net/science/chem/exps/Ni_en_comple...

blargish - 8-10-2014 at 15:39

Quote: Originally posted by No Tears Only Dreams Now  


No no no, I've spent a lot of time doing copper chemistry, copper(II) hydroxide is insoluble, it makes for cloudy, gross colored slurries. Also, this solution turns violet(blue+magenta) when phenol red indicator is added, indicated that it's strongly basic. Here's the solution on its own:



I know that copper hydroxide is insoluble. I was just wondering if maybe some minute suspension of it was formed in the solution. Is your solution completely clear? Or cloudy? I can't entirely tell from the image

The solution would certainly be basic due to your adding of the excess bicarbonate and hydroxide; although I'm not sure exactly what is counting for the color change upon the addition of NaOH.

Amos - 8-10-2014 at 15:44

Quote: Originally posted by blargish  
Quote: Originally posted by No Tears Only Dreams Now  


No no no, I've spent a lot of time doing copper chemistry, copper(II) hydroxide is insoluble, it makes for cloudy, gross colored slurries. Also, this solution turns violet(blue+magenta) when phenol red indicator is added, indicated that it's strongly basic. Here's the solution on its own:



I know that copper hydroxide is insoluble. I was just wondering if maybe some minute suspension of it was formed in the solution. Is your solution completely clear? Or cloudy? I can't entirely tell from the image

The solution would certainly be basic due to your adding of the excess bicarbonate and hydroxide; although I'm not sure exactly what is counting for the color change upon the addition of NaOH.


I would guess that there may be complexation of the copper ions in solution; that would explain why they didn't form a precipitate, and if the complexation was favored by strongly basic conditions, that might account for the coloration.

CuReUS - 9-10-2014 at 08:41

could choric acid be used to replace nitric acid in aqua regia?

because in aqua regia ,nitric acid does the same work of oxidising the surface of gold:o

DraconicAcid - 9-10-2014 at 08:48

Quote: Originally posted by Brain&Force  
It probably has something to do with the presence of chlorate. It reminds me of this:

http://woelen.homescience.net/science/chem/exps/Ni_en_comple...


I know Woelen had some post about iodate or periodate complexes of copper- I don't know if chlorate is much of a ligand. Perchlorate, at least, is quite non-coordinating.

Brain&Force - 9-10-2014 at 09:57

Maybe you've made a cuprate complex? blogfast25 says it's easy to demonstrate with amateur equipment.

CuReUS, I think it would decompose to form nasty chlorine and chlorine oxides.

Little_Ghost_again - 9-10-2014 at 10:38

Take some and make a crystal, would be cool to see the shape

Amos - 9-10-2014 at 11:07

Quote: Originally posted by Little_Ghost_again  
Take some and make a crystal, would be cool to see the shape


At this point I'm dealing with nearly 200mL of solution, which contains AT MOST half a gram of copper; there's a whole lot more sodium carbonate and sodium hydroxide in there, but maybe I can reproduce this without basification.

Amos - 9-10-2014 at 11:10

Quote: Originally posted by CuReUS  
could choric acid be used to replace nitric acid in aqua regia?

because in aqua regia ,nitric acid does the same work of oxidising the surface of gold:o


Y'know, I think it would, being so powerfully oxidizing and being about as strong as nitric acid, but I imagine it would be much more dangerous, as you'd be working with concentrated chloric acid which is VERY unstable, and the reaction is exothermic, which chloric acid also hates.

CuReUS - 11-10-2014 at 05:45

Quote:
but I imagine it would be much more dangerous, as you'd be working with concentrated chloric acid which is VERY unstable, and the reaction is exothermic, which chloric acid also hates.


but isnt copper much more reactive than gold
also i thought that since you said that no chlorine gas or anything was released it would be good because nitric acid in aqua regia spewes out nitrogen dioxide,nitrosyl chloride:mad:

blargish - 11-10-2014 at 19:37

Quote: Originally posted by CuReUS  
Quote:
but I imagine it would be much more dangerous, as you'd be working with concentrated chloric acid which is VERY unstable, and the reaction is exothermic, which chloric acid also hates.


but isnt copper much more reactive than gold
also i thought that since you said that no chlorine gas or anything was released it would be good because nitric acid in aqua regia spewes out nitrogen dioxide,nitrosyl chloride:mad:


Chloric acid would no doubt consistently emit chlorine oxides, especially at higher concentrations. Even so, I'm not even sure how viable this idea is since you really cannot obtain a concentrated chloric acid solution due to its instability...