The Long Road to Copper (I) Iodide

It's been a while since my last post, because I've focused more on my YouTube account (linked on the main page). However, recently I did a rather long chain of experiments that I haven't taken video of, and so I thought would work better as a blog post with a few photos. This post, as the title suggests, is long. You were warned.

This experiment was a great example of "real science." I call it that not to imply that the rest of my stuff is somehow fake, but rather because this whole process was new to me and while I had a general idea of what to do I wasn't sure what to expect. Some of my other posts and videos are more along the lines of a "demonstration" than an actual experiment. This one was fun in that a number of problems and strange situations cropped up, and I had to think up ways to get around them. Anyways, on to the meat of the post.

This all started with the flame coloration video that I've been working on. Basically, if you burn different metal salts, they color the flames different colors. So far I am able to create all the colors of the rainbow except blue. That requires a copper (I) compound (other things can work, but that's what I chose to shoot for). I've reached the point in my chemistry hobbyist career where I have enough chemical stock to make most things I need rather than buying them, which is a lot more fun and I don't have to wait for shipping. So, to sum things up, I needed to make copper (I) iodide. The reactions involved are the following:

2KI + CuSO4 ==> CuI2 + K2SO4

2CuI2 ==> 2CuI + I2

Copper (II) iodide is unstable, and rapidly dissociates into copper (I) iodide and iodine. I have copper (II)sulfate (CuSO4, sold as stump remover), but didn't have potassium iodide (KI). So the first step was to make that from things I did have, via the reaction between potassium hydroxide and elemental iodine.

6KOH + 3I2 ==> 5KI + KIO3 + 3H2O

The potassium hydroxide solution must be hot, and luckily this is automatically taken care of because dissolving KOH flakes in water generates a lot of heat. You can see the iodine reacting with the KOH solution in the photo below. I actually had to heat it a bit extra to get everything into solution.

It's very easy to tell if you've put too much iodine in, because the solution immediately goes yellow or brown. Iodine is not very soluble in water, but becomes much more so in the presence of the iodide ion by forming the triiodide ion I3-. This is easily solvable by adding a few chips of potassium hydroxide until the solution becomes clear again.

The coproduct of potassium iodate (KIO3) is easily separated by the difference in solubility. KIO3 is much less soluble in water than KI, and most of it will precipitate out as a solid. To remove as much of this as I could, I further placed the beaker in an ice bath and got a bit more precipitate. My iodate sample was tan because of iodine impurities, which I can remove by the same process I'll use later down in this post. 

So, now I have my solution of potassium iodide that I need for the main reaction. Upon mixing with the deep blue solution of copper (II) sulfate the color immediately changed to a very dark brown, looking like a cloud of mud. I then filtered this mixture, and this is where things got weird. I actually came up with a green solid.  

 
Copper (I) iodide is white when pure, and mine would likely be tan like the potassium iodate above. The remaining solution was a very dark brown, because of the free iodine released in the secondary reaction. So at least that part of it worked. I tried to use this green solid for the flame coloration experiment, but got no results (it didn't show any color at all).

At this point, I wasn't sure what was going on so I left everything alone for a few weeks while I tried to decide what to do. I didn't want to waste the iodine in solution, because there should be a way to recover it.

After some time had passed, I revisited this experiment once more. About half of the water had evaporated from the solution, and I heard some solids clinking around at the bottom of the beaker. Expecting this to be crystallized iodine, I filtered it again. Once more, I was surprised to discover that this time green crystals had appeared.

This must be the same solid as before, just in crystal form because it had much longer to leave solution. I've since come to suspect that this must be copper carbonate, the same material that makes old copper items green (essentially copper rust). There's a test I can perform to check this hypothesis: By heating it to 290 C, copper carbonate will decompose into black copper (II) oxide and carbon dioxide gas. I haven't tried this yet, but in any case I've got no idea where the carbonate could have come from. I'd say impurities in my chemicals, but I think I produced way too much of it to attribute it to that.

At this point, I decided to repeat the main experiment again just to see what would happen. I made another solution of copper sulfate and added it until the reaction appeared to slow down significantly. This time upon filtering, I recovered a purplish solid! I believe this is the copper (I) iodide I was looking for, just heavily contaminated with iodine.

To remove the iodine, I washed the solid with an ice-cold solution of sodium thiosulfate, Na2S2O3. Sodium thiosulfate is capable of reducing elemental iodine to iodide, which is harmless and soluble in water. The water is chilled to prevent dissolving and carrying away any of the copper (I) iodide. After performing this process three or four times, my copper (I) iodide has become a much lighter tan color and is much more pure.

The leftover iodine solution remains as dark as ever, and I'd love to find a way to transform it back into iodine crystals. I've tried a chemical method of mixing with hydrogen peroxide, but this didn't yield anything. Probably the best way is to simply distill it and condense the iodine vapor that comes out. If all else fails, I can always dispose of it by reduction with sodium thiosulfate. But, I'll leave that for another day.

In the end, I tried the flame test once more with this new product. Once again, it failed to color the flame. So after all that, it still didn't work! But I don't view any of this as a failure. In chemistry, there is no such thing as failure. Anything that doesn't work as you expect it to becomes a learning experience, and provides great information and experience for anything you try to do in the future. It's also a lot of fun to me, to be able to explore all the different, sometimes surprising, things that can happen in the chemistry lab, no matter what the outcome.

Well, I hope you enjoyed reading this, and congratulations for making it to the end! Remember to check out my YouTube page for more frequent updates, and more chemistry!

If you want to learn more about the element iodine, ChemTalk has a good article on the subject: Intriguing Iodine. Check it out to learn some cool facts and read about some other interesting chemistry that's possible with this versatile element.