Make SULFURIC acid from WEAK oxalic acid with the SUN

Greetings fellow nerds. Every couple of years or so I show another way to make sulfuric acid, it seems to have become a tradition at this point. In this video we're going to do it by a rather unusual process, oxalic acid precipitation. I'll tell you upfront the yield was around 46% and by overloading the reagents as well as using unorthodox reducing agents we can bump that yield up to 67%. So amateur viable, but since it uses a lot of ferrous sulfate and oxalic acid, it’s very expensive compared to the sulfuric acid produced. So let's get started. First we get 300g of iron (II) sulfate heptahydrate, also called ferrous sulfate heptahydrate. I bought mine as a gardening chemical. Then we dissolve that in 1L of water. You may need to gently heat it to speed up the dissolution. While that is going we get exactly half the mass of oxalic acid dihydrate, or 150g and dissolve that in 300mL of water. We heat the oxalic acid to dissolve it. Then once it's dissolved we add that to the previous solution of iron sulfate. Oxalic acid has the rather interesting property that it forms insoluble oxalate salts with many metal ions. So if reacted with iron sulfate it will form insoluble iron oxalate. The sulfate ion is freed and becomes sulfuric acid. This is a rather unusual reaction because usually the reverse reaction is more favorable. Oxalic acid is a much weaker acid than sulfuric acid and usually if you react a weak acid salt with a strong acid you make the weak acid and the strong acid salt. If you react a strong acid salt with a weak acid, like we are here, usually nothing happens. But in our case, the iron oxalate is so insoluble that it drives the reaction forward and we create the stronger sulfuric acid despite starting with a weaker acid. This doesn’t work with other weak acids like acetic acid. So oxalic acid is quite unusual like that. Anyway, keep stirring until it cools and then let it settle. This is so that it's easier for us to decant and filter it later. I find it much easier to decant and filter rather than filtering the slurry directly. The residue is iron oxalate and is useful in itself for making such interesting compounds like pyrophoric iron, so you can save it if you want. Although it is very impure containing lots of unreacted oxalic acid, so keep that in mind. Meanwhile our sulfuric acid in the filtrate and i boiled it down until it started to crystalize. For my solution this was around 200mL. Then we let it cool. The unreacted iron sulfate is much less soluble in sulfuric acid solution so it crystalizes out. We then decant that into a new flask and here you can see the iron sulfate. We needed to remove this so it won’t bump when we distill the sulfuric acid. So now we setup a sulfuric acid distillation system and distilled our sulfuric acid. For further information on how to distill sulfuric acid refer to my previous video on sulfuric acid purification, links are in the video description. I will add though that boiling oxalic acid with sulfuric acid produces small amounts of carbon monoxide, so this distillation must be performed outside or in a fume hood. You should still be doing that anyway when distilling sulfuric acid, but I thought I should mention that. I kept checking the density of the distillate as it came off and when the density was above that of water, I started collecting the sulfuric acid. And there it is, about 70 grams of diluted sulfuric acid. To find the concentration we do a density check and i got 1. 61g which gives us a concentration of about 70% for this temperature. So our 70 grams of dilute acid contains 49g of actual sulfuric acid. The expected yield was about 105. 7g so our yield was 46%. So the yield isn’t terrible, but considering we had to process 450g worth of chemicals just to get a tiny amount of acid i don’t think it’s viable overall. But i’ll leave that choice up to you. The yield isn’t high primarily because the reaction doesn’t go to completion. As said before reacting a strong acid with a weak acid salt tends to produce the weak acid and strong acid salt. So as sulfuric acid builds up, it reacts with the iron oxalate and converts it back to iron sulfate and oxalic acid. You can push the reaction forward somewhat by diluting the chemicals with even more water to reduce the overall concentration of sulfuric acid. I couldn’t do that because i was limited by my glassware.

But one thing we can do is use a much greater excess of reagents. So i tried again but this time i used 200g of oxalic acid rather than 150g. This much greater excess of oxalic acid pushes the reaction forward again and after doing the subsequent filtering and distilling i was able to get about 57% yield of sulfuric acid with the displayed parameters. So this process can be optimized. Another contributor to the reduced yield was the oxidation of iron 2+ ions to iron 3+ by oxygen in the air. You can clearly see it as this yellow brown color of the iron sulfate when pure iron 2+ sulfate should be green. The resulting iron 3+ oxalate is somewhat more soluble, especially under these acidic conditions. Some amateurs improve their yield by adding a reducing agent like ascorbic acid or bubbling in sulfur dioxide. Perhaps i’ll try it in the future. But instead i’m going to redo the experiment and try a different reducing agent. I’m going to use a nice summer’s day and the full concentrated power of the sun. i’m not joking, I've redone the experiment with the exact same proportions as earlier with the 200g of oxalic acid and before I filtered the iron oxalate I left the beaker in direct sunlight for 8 hours. What’s happening is that iron (iii) oxalate is actually photosensitive and photo decomposes into iron (II) oxalate and carbon dioxide. This is actually a well known reaction and used for some types of early photography before the development of more advanced silver halide based processes. A great benefit for us is that it’s clean and the resulting carbon dioxide easily boils out. Anyway, after 8 hours in the sun, I brought it inside, filtered it, boiled it down, distilled it and was able to obtain a 67% yield. I’m actually impressed that simply leaving it in the sun created such a big yield improvement. Maybe if i left it out for a week i’d get even better yield. But i’ll leave those experiments to someone else, I’m too lazy. So i hope that video was useful to you. Before i go, i’m also going to look at a similar process of using magnesium sulfate rather than iron sulfate. I’ll tell you upfront it failed and only gave 2% yield so feel free to stop the video now, or watch me analyze why it failed. Magnesium sulfate and oxalic acid is talked about a lot in online chemistry message boards and sounds like a cheaper alternative to iron sulfate. But there are conflicting reports on its viability. So I made my own experiment using 500g of magnesium sulfate heptahydrate in a liter of water and 300g of oxalic acid in 500mL. After boiling them and mixing them I got a precipitate and I thought it was working. So I filtered them, boiled them down and then proceeded to distill the acid. The final amount I obtained was 25g at a density of 1. 1g for a concentration of about 15%, or about 3. 75g. I was expecting 200g or so and the yield was under 2%. 2% is completely useless for the amateur. You might as well spend the same money buying copper sulfate and electrolyzing it, the previously most expensive way to make sulfuric acid. At least this magnesium oxalate process now gains that title. It seems this process failed because the magnesium oxalate formed just wasn’t stable enough to drive the reaction forward. So the sulfuric acid produced simply reacted with the magnesium oxalate and drove it back again. But I can understand why so many amateur chemists were fooled, they weren’t lying, they just misinterpreted their results. If you look at the distillation you can see a massive amount of magnesium sulfate residue in the boiling flask. What I think happened was amateur chemists that tried this measured the density of their filtrate and found it was much higher than water and so assumed it was sulfuric acid, not realizing it was just dissolved magnesium sulfate. This belief suffered further confirmation bias when they saw the precipitation during the initial mixing. This precipitate contains some magnesium oxalate, but actually its mostly oxalic acid cooling and crystalizing back out. Anyway, the small amount of sulfuric acid that was produced as well as dissolved oxalic acid gave an acid indication to pH testing. So the bias was confirmed and those chemists declared this method successful. I don't think they were, but I don’t think they were lying. Distilling sulfuric acid to test it is extremely dangerous and difficult and I don't blame them for not trying it. It’s really hard to be an amaetur chemist and I myself have made countless mistakes, biased conclusions, and supremely ignorant theories. Finally i want to explore the use of copper sulfate.

is also pretty easy to get although it’s somewhat more expensive than iron sulfate. Nonetheless it also forms insoluble oxalates so it’s worth trying out. There are some reports on chemistry message boards claiming this was superior to iron sulfate. So I prepared a mixture consisting of 250g of copper sulfate and 150g of oxalic acid in a total of 1. 5 liters of water. Interestingly enough it did work and it seemed to produce a proper precipitate of copper oxalate. Unfortunately when i tried to filter it, the copper oxalate was so fine that it passed right through. It’s like trying to filter milk, although this looks closer to star wars blue milk that Luke Skywalker drinks. Anyway, even letting it sit for a day to settle didn’t help much. I also tried boiling the mixture to induce ostwald ripening but that didn’t work either. I added excess copper sulfate because maybe you need excess copper to get the particles to stick. Finally I tried adding 10g of iron sulfate to act as a flocculating agent but that also turned up nothing. This is disappointing because the process does seem to work. The minor amount of solution that separated was almost clear, indicating the copper ions were indeed mostly removed. But without a way to separate the copper oxalate it doesn’t seem viable. Nonetheless i’m probably doing something wrong, various other amateurs report success so maybe there is some small difference in impurities between their chemicals and mine that allows for separation. But overall, i would still use the electrolytic method to convert copper sulfate into sulfuric acid since that’s cheaper than oxalic acid. So thanks for watching. If you like the video please subscribe, like and comment. I know some of you will ask if this method can work for nitric acid. Actually it does, and I'm currently working on a video for that. So please subscribe to stay informed. If you really like my content, then consider becoming a patreon member where you get video previews of works in progress. Alternatively, one time donations with youtube superthanks is greatly appreciated.