Again, very nice, thank-you. I've been playing around, somewhat, trying to understand the Copperas method of making sulphuric acid from iron pyrites (iron II disulphide). Wiyhout, so far very much success. However, I've found that aqueous solutions of iron II suphate seem to oxidise rather quickly at room temperature to a rather gruesome red, then yellowish muck. Incidentally, the 'Ferrous Sulphate' that I've had as a box of garden fertilzer Ofor some time, unopened, but in a cardboard box, seems already to have oxidised in the same way to produce an insoluble yellow/brown solid. Hmmm . . . Aggain, very many thanks for your excellent work!
This was an excelent tutorial! I found Your channel milion times more interesting and informative than about 99% of those hundredthousand subs channels out there.
Good video, been looking a step by step video like this one. Great job. I need this for gold precipitation so this is a great find! Thanks. Will be sharing this with my hobby refiners if you don't mind. :)
Go ahead and share this with anyone you wish. The Mohr's salt is stable and stores well over time but the ferrous sulfate tends to oxidize and rust rather easily. I have some ferrous sulfate that I have kept for a few months now and it is doing fine. The trick was to fill the container with CO2 before I sealed it up airtight.
@ChemicalMaster Dry ice would work but I did it by making a simple CO2 generator using baking soda and vinegar. I directed the gas into the container with a piece of tubing after covering the top of the container with some plastic wrap. The CO2 is heavier so it will displace the air inside the container. You have to guess when you think it is full then simply remove the hose and plastic wrap. Install the lid and you are done.
Im curious if you know i also used this exact steel wool for my iron compounds and im curious what the impurities are if you know or guess on. Im guessing chromium is a prominent one but I could be wrong.
The weight was close to the theoretical ammount for the heptahydrate . The color matches the heptahydrate. . The temperature would have to be at or above about 55C while it dries to form the lesser hydrate forms . I allowed the drying to occur at room temperature.
You have a nice professional surface white tiles are excellent.It did not look like you had the iron sulphate heptahydrate. It looked very white. Too little water was used. There are a range of iron sulphate salts. If a solution is boiled and the moment the crystals form, turn off the hotplate and cool very slowly with cloth wrapped around it. Then you get the nice blue green crystals. There is also a solubility table on wikipedia which would help avoiding those hiccups.
@ARTHUR Broquet There are a couple of reasons why I think it was heptahydrate. There was excess water when it was air dried at room temperature so the heptahydrate should have formed naturally.. The mass was very close to the expected amount if heptahydrate was assumed. The color could appear more white than expected because the crystals were very small and diffract light differently. If you take a larger blue green crystal and grind it to a powder it will appear more white.
@@chemicalmaster3267 Until you mentioned it I didn't even know ferric ammonium sulfate was a thing. I'll keep it on my list of possible future experiments. I have several thing I want to do first though so don't hold your breath. :) I did find a way to prevent my ferrous sulfate from oxidizing and turning brown. One time I made a batch and had some beautiful crystals. I placed them in an airtight plastic bottle for storage. I noticed after a few weeks a lot of them had turned a brown rusty color and the sides of the plastic bottle were sunken in. All of the oxygen had been consumed and created a partial vacuum. After I made a fresh batch I then filled the bottle with carbon dioxide and sealed it tightly. After a couple of months they are still looking good.
@@perrygershin3946 Yeh, most ferrous componds are quite easily oxidized by oxygen in the air. For now I can only think of one iron(II) compound that is immune to oxidation by oxygen: ferrous oxalate dihydrate.
@@perrygershin3946 Also what do you think, would you consider doing a video about the synthesis of potassium ferrioxalate trihydrate? It forms beautiful emerald green crystals but be careful to not leave its solution exposed to UVs for too long or it will decompose. All that it´s needed is potassium hydroxide or carbonate, ferric(III) oxide or hydroxide and oxalic acid.
Again, very nice, thank-you.
I've been playing around, somewhat, trying to understand the Copperas method of making sulphuric acid from iron pyrites (iron II disulphide). Wiyhout, so far very much success.
However, I've found that aqueous solutions of iron II suphate seem to oxidise rather quickly at room temperature to a rather gruesome red, then yellowish muck.
Incidentally, the 'Ferrous Sulphate' that I've had as a box of garden fertilzer Ofor some time, unopened, but in a cardboard box, seems already to have oxidised in the same way to produce an insoluble yellow/brown solid.
Hmmm . . .
Aggain, very many thanks for your excellent work!
This was an excelent tutorial! I found Your channel milion times more interesting and informative than about 99% of those hundredthousand subs channels out there.
Good video, been looking a step by step video like this one. Great job. I need this for gold precipitation so this is a great find! Thanks. Will be sharing this with my hobby refiners if you don't mind. :)
Go ahead and share this with anyone you wish.
The Mohr's salt is stable and stores well over time but the ferrous sulfate tends to oxidize and rust rather easily.
I have some ferrous sulfate that I have kept for a few months now and it is doing fine.
The trick was to fill the container with CO2 before I sealed it up airtight.
@@perrygershin3946 How did you do that? Dropped dry ice inside the container?
@ChemicalMaster
Dry ice would work but I did it by making a simple CO2 generator using baking soda and vinegar.
I directed the gas into the container with a piece of tubing after covering the top of the container with some plastic wrap.
The CO2 is heavier so it will displace the air inside the container.
You have to guess when you think it is full then simply remove the hose and plastic wrap. Install the lid and you are done.
You can use oxidation of Mohr’s salt to measure ionizing radiation exposure, particularly gamma.
Im curious if you know i also used this exact steel wool for my iron compounds and im curious what the impurities are if you know or guess on. Im guessing chromium is a prominent one but I could be wrong.
Your channel is great
Keep up with the good work 👍👍👍👍👍
How to make copper sulphate pentahydrate.?
You did great. I am trying to do ink (with tanin). I hear that ink was really acidic and is eating the paper.
Would baking powder help?
How do you know iron sulfate is in heptahydrate form ?
The weight was close to the theoretical ammount for the heptahydrate
.
The color matches the heptahydrate.
.
The temperature would have to be at or above about 55C while it dries to form the lesser hydrate forms
.
I allowed the drying to occur at room temperature.
@@perrygershin3946 Ok , thanks Perry , i suppose i can follow the same procedure to make this product .
Iron ii charge with sulphuric di edtnol. P j great reyan carbide acidii felore mehuir solu 2sulphite one hydolite. Gives ?
You have a nice professional surface white tiles are excellent.It did not look like you had the iron sulphate heptahydrate. It looked very white. Too little water was used. There are a range of iron sulphate salts.
If a solution is boiled and the moment the crystals form, turn off the hotplate and cool very slowly with cloth wrapped around it. Then you get the nice blue green crystals. There is also a solubility table on wikipedia which would help avoiding those hiccups.
@ARTHUR Broquet
There are a couple of reasons why I think it was heptahydrate.
There was excess water when it was air dried at room temperature so the heptahydrate should have formed naturally..
The mass was very close to the expected amount if heptahydrate was assumed.
The color could appear more white than expected because the crystals were very small and diffract light differently.
If you take a larger blue green crystal and grind it to a powder it will appear more white.
+Perry Gershin Would you do a video where you make ferric sulfate and ferric alum (ferric ammonium sulfate)?
I'm sorry but I do not understand this question.
@@perrygershin3946 Whoops! I forgot to add some words. I corrected the question right away.
@@chemicalmaster3267
Until you mentioned it I didn't even know ferric ammonium sulfate was a thing. I'll keep it on my list of possible future experiments. I have several thing I want to do first though so don't hold your breath. :)
I did find a way to prevent my ferrous sulfate from oxidizing and turning brown. One time I made a batch and had some beautiful crystals. I placed them in an airtight plastic bottle for storage. I noticed after a few weeks a lot of them had turned a brown rusty color and the sides of the plastic bottle were sunken in. All of the oxygen had been consumed and created a partial vacuum. After I made a fresh batch I then filled the bottle with carbon dioxide and sealed it tightly. After a couple of months they are still looking good.
@@perrygershin3946 Yeh, most ferrous componds are quite easily oxidized by oxygen in the air. For now I can only think of one iron(II) compound that is immune to oxidation by oxygen: ferrous oxalate dihydrate.
@@perrygershin3946 Also what do you think, would you consider doing a video about the synthesis of potassium ferrioxalate trihydrate? It forms beautiful emerald green crystals but be careful to not leave its solution exposed to UVs for too long or it will decompose. All that it´s needed is potassium hydroxide or carbonate, ferric(III) oxide or hydroxide and oxalic acid.
Super sir
Thank you.
Sulphurtroibine nonmono fatty soil. Fessophite baruin kilopotaste netrate one hydolite calbide triem. Carb hydrate sul.monochrome.
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