Cool Experiments and Techniques page

[mevans]

Im starting this thread to post any ideas or techniques i have that may help you guys out. Anyone who wishes to should feel free to comment or write add their own experiment to this page.
I will start off with one on how to purify Fe3O4
if any of you guys are like me, then you probably have some Fe3O4 in your lab for thermite. The problem is that the reagent grade stuff is expensive. Ungraded Fe3O4 is avalible for purchase but it can be highly impure. This annoyed me since i also wanted to use it for the creation of iron salts. Because it is not soluble in water, you cannot do purification by recrystalization and since most of the impurities are insoluble as well, you cannot filter it. Fortunately there is another way around this and in no time you can turn a 50+ gram ungraded sample of Fe3O4 into a much more pure product.

*This process will not work with Fe2O3 as it is not magnetic

Step 1- Weigh out 50 grams of Fe3O4 (If you want you can write down its starting weight and subtract it from its final weight and calculate the original purity.)
Step 2- Put the Fe3O4 in a 250ml beaker and add about 200ml or less of tap water. Then stir the contents.
Step 3- if you bought the $2.75 a pound stuff like i did then the impurities will be pretty obvious when you start stirring.
Step 4- Stir for about 5 seconds and then get a bar magnet and put it on the bottom of the beaker. Since Fe3O4 is magnetic, most of it should immeadiately settle at the bottom of the beaker.
Step 5- Put a funnel in a large container and then pour off the dirty water into the container.
Step 6-Repeat steps 2-5 untill the water is almost clear.
Step 7- do at least 2 more rinses but this time use distilled water
Step 8- after your final rinse add some acetone and then pour it out of the beaker (this step is optional and just speeds up the drying)
Step 9- put the purified Fe3O4 in an evaporating dish untill dry (if you wrote down the starting weight then you can now determine original purity.)
Step 10- depending on how many decantings you did, your product should be sufficiently pure for making iron salts ect.

[mevans]

Here is a second one that is a cool experiment that anyone with a basic lab can do at home.
First off let me say that this was not my idea i mearly found it while browsing instructables.com.
In this experiment post i will be describing how to isolate the chemical compound in splenda that makes it sweet called 1,6-Dichloro-1,6-dideoxy-β-D-fructofuranosyl-4-chloro-4-deoxy-α-D-galactopyranoside , (aka sucralose). Sucralose is (please correct me if im wrong. ) a molecule of sucrose that has been chlorinated. The resulting product is a compound 500 times sweeter that table sugar. "Then why dosent splenda taste any sweeter than sugar" some of you may be thinking. The answer is becuase most of splenda is a tasteless substance known as malodextrin with a small portion of sucralose mixed in. Fortunately for us we can isolate pure sucralsoe with an organic solvent such as 2-propanol (isopropyl alcohol), 2-propanone (acetone) or ethanol (ethyl alcohol). This will dissolve the sucralsoe while leaving the malodextrin in solid form.

a lewis diagram of sucrose

Sucrose_CASCC.png (5.52 KiB) Viewed 250 times

a lewis diagram of sucralsoe

1000px-Sucralose2_svg.png (41.42 KiB) Viewed 250 times

Step 1- Get some splenda. ( i got about 2 cups worth of it) and put it in a beaker.
Step 2- Add some of your organic solvent and stir for 5+ minutes.
Step 3- Filter out the malodextrin that did not disolve.
Step 4- Pour your solution into an evaporating dish and let it evaporate on its own.
Step 5- once it begins to evaporate add more of your sucralose-solvent mixture to the evaporating dish untill you run out
After all of the solvent has evaporated (it took about a week with 2-propanol as the solvent) I was left with a few crystals of sucralsoe. Just sniffing it made my mouth water. Im sure the question on your mind is can i taste it. I would have to tell you no simply becuase i do not want to be responsible for injury. I did taste it and it was quite an experience. I would say absolutely not if you used acetone. Any questions or comments please post
Thanks,
Mevans]

[mevans]

A way to identify the concentration of copper compounds in a solution

I came up with this idea when i observed the reaction between potassium iodide and copper sulfate. The way this works is similar to a regular titration. In a flask, place the solution that you want to determine the copper concentration of. In the burette, add a standardized solution of .1 M potassium iodide. As you add the KI, you will notice the formation of copper(I) iodide which is formed because of the excess of copper ions.
2I- + Cu2+ --> CuI
CuI is insoluble and forms a precipitate on the bottom of the flask. Eventually, the Cu ion will no longer be in excess and the KI will. This will begin to change the insoluble CuI into soluble copper(II) iodide (CuI2).
CuI + I- ---> CuI2
CuI2 instantaniously decomposes into Copper(I) iodide and elemental iodine.
2CuI2 --> 2CuI + I2
The iodine turns the solution a purple/brown color and if the solution is allowed to settle, the dissolved iodine will float to the top.
Somthing to note is the that when a solution gets close to the titration endpoint, the dissolved elemental iodine will appear in solution and eventually dissapear once it reacts with copper ions left in solution.
This technique gave me promising results when i used it on a copper sulfate solution of a known concentration (the solution was 1.00M and the test said it was 1.03M). I belive that this technique would also work with other oxidation states of copper and mabe in qualitative and quantitative analysis.
Well I hope this helps. If you do use this technique, please comment and share your experiences. I am always open to suggestions, criticisms, and complements so make sure to speak up.
Mevans