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Everything posted by Sebbass69

  1. Hi everyone I have a quick question I hope you all can answer. I'm titrating some HNO3 to determine concentration, however I have not standardized my NaOH solution. I'm looking for a suitable standardizing agent, and I've come across a few. KHP seems to be the most common but oxalic acid and hydrochloric acid are also listed. If I was to buy a standardized solution of HCl, would that be suitable for standardizing my NaOH? I can buy standardized NaOH but doesn't it frequently absorb moisture and CO2 from the air thereby altering concentration and throwing off the titration? Thank you, ~Sebastian
  2. According to the Wikipedia article, bromine is shipped in steel containers that are lined with lead. Perhaps the lead is unreactive enough to resist corrosion by the bromine? Or perhaps it is just easily passivated... anyone have any thoughts?
  3. Hey guys, So, this is going to sound very strange but bear with me... I was wondering if anyone knows of a way to remove the graduations and labeling off of glassware without scratching or destroying the glass. Just because this is such a strange request, I'll explain: It's actually just because I'm so anal about aesthetics I'm thinking of buying some media bottles that have graduations and labels and stuff, and I don't like the way the markings look, so I wanted to remove them. I know that's weird, but it's just how I am... sigh Anyways, thanks for your help! ~Sebastian P.S. Sorry if this is in the wrong section, but I wasn't sure where to put it...
  4. I think you may be correct in saying that the amount of chlorine generated is not of real concern, however, chlorine bleach is definitely more reactive, and more able to generate chlorine gas than table salt is. I wouldn't be so scared of mixing concentrated HCl with table salt, but with bleach that would be a whole different story. That and I would challenge you to eat a teaspoon of table salt, and then a teaspoon of bleach. Saying the two are equivalent is just not correct....
  5. I feel like it would just be so much easier (and safer) to buy it....
  6. I've been reading about toluene, and I understand the distinction between dry and wet toluene. However, I was wondering how much water toluene actually absorbs, and if it is really important to have dry toluene for general chemistry. If it is absolutely critical what is the easiest way to dry toluene?
  7. I have 35% Food Grade H2O2 on hand, and I'm tired of worrying about it decomposing, but buying stabilized H2O2 is very expensive. I then thought to myself, well, why don't I just stabilize it myself. So, after some searching through Google, I found that the most common stabilizers are Acetanilide (used in USP H2O2), Sodium stannate, phenol (which I don't want to use because it causes cancer, if I recall correctly) and Tetrasodium phosphate (AKA Sodium Pyrophosphate). I also found this :"Colloidal stannate and sodium pyrophosphate (present at 25 - 250 mg/L)" - which gives the approximate quantity of stabilizer per liter, for those two substances. As well, I'm going to assume that colloidal stannate and sodium stannate are interchangeable. Now that I have written you all an essay, my question is: Does anyone know which stabilizer is most effective? Would it just be best to add all three to the peroxide? And what would be an appropriate amount of acetanilide to add, as I couldn't find a number for that?
  8. So I was at the pharmacy the other day, and noticed they had mineral oil and needing some to store sodium, I looked at the purity and saw that it was 99.9% pure. However, I had also found some from a chemical supplier, and I wasn't sure if theirs would be better - so I figured I would poll the audience- does everyone think / know if the mineral oil from the pharmacy would be pure enough to store my sodium, or should I get it from a chemical supplier? I try and stay away from buying things OTC (like H2SO4 drain opener, due to heavy impurities) but this is from a pharmacy, and is 99.9% so it seems like it would be pure enough...
  9. Epe (pronounced Epay) is an extraordinarily hard wood - its like stone - my dad worked on a house where they had to cut it to build a small porch, and he was telling me how he had to replaced the blade on his table saw twice because the stuff is so hard only a fresh blade can cut it - Only problem is that its really endangered, so its use is kind of frowned upon - however, to make a sword I imagine you would use so little, it wouldn't matter.
  10. why would you do a titration of anhydrous ammonia? Anhydrous by definition means that it is free of water, and therefore must be 100%... Or am I misunderstanding your experiment?
  11. Yeah, I got that info from Wikipedia - "Ammonium persulfate was prepared by H. Marshall by the method used for the preparation of potassium persulfate — by the electrolysis of a solution of ammonium sulfate and sulfuric acid." - Hugh Marshall (1891). "LXXIV. Contributions from the Chemical Laboratory of the University of Edinburgh. No. V. The persulphates". J. Chem. Soc., Trans. 59: 771. doi:10.1039/CT8915900771
  12. So, I know that acids are very conductive of electricity. However, what kinds of materials are liberated when you run a current through an acid? For example, with HCl, I imagine it would be hydrogen and chlorine, but what about sulfuric acid? How about phosphoric? I know that when ammonium sulfate is added to sulfuric acid and then electrolyzed, it forms ammonium persulfate, but what about without any substance in the acid?
  13. I performed this reaction, and it worked perfectly. However, I don't seem to be able to dry the material I have produced. Normally, I would heat it in a small oven, but the Ba(OH)2 is reactive with carbon dioxide, and therefore would react while drying. My first though was to make a chamber, with a small dish of sodium hydroxide, and a small dish of barium hydroxide to be dried. The Ba(OH)2 would be sealed, and the sodium hydroxide would be open. The NaOH would absorb all the CO2 from the chamber, and after all the CO2 is absorbed, the dish with the barium would be opened, and the evaporated water would be absorbed by the NaOH remaining. If need be, a small heating element could be added above or below the Ba(OH)2 to drive off any water. The only other way I could do this, would be to make a completely sealed oven, and perform the same process with NaOH to remove CO2. Anyone have any thoughts, or simple ways of drying the barium hydroxide?
  14. Whether or not you believe me is irrelevant - you're belief or disbelief does not make it true or untrue - however I assure you, I work in a school lab - Merged post follows: Consecutive posts mergedI forgot to say - from what everyone tells me, it seems that actually concentrating H2O2 isn't really feasible without specialized equipment, so I think it's safe to say it would be foolish to attempt it - my school is good, but they don't have vacuum distillation apparatus that can handle H2O2 - that and from what you guys say, it's explosive, so I think it would be best to wait on this one - Anyways, thanks for all of the advice guys -
  15. Fair enough - but you're still forgetting that I'm going to be working with my teacher, who is a professional, and would provide the required equipment - You are right about one thing - I was asking about it's properties, so I would have some idea of how to do this.... Additionally, no MSDS on 90% H2O2 lists it as an explosive - it says "Product is non-combustible. On decomposition releases oxygen which may intensify fire. An explosion hazard when mixed with organics at high concentrations." That's why I thought it wouldn't explode by itself unless heated - I knew it could explode when mixed with flammable / combustible material, but if you have other experience, then thank you for the information - As for your remark, it could have been more polite - wisdom is not a substitute for manners, nor is it an excuse to be rude.
  16. thank you for your help - Yeah, I know that water soluble barium salts are very toxic -
  17. Would it be possible to synthesize Ba(OH)2 from an aqueous solution of sodium hydroxide and barium chloride. The reaction would look something like: BaCl2 + 2 NaOH --> Ba(OH)2 + 2 NaCl Because Ba(OH)2 is not nearly as soluble as the other salts, it would fall out of solution - not all of it, but some of it. Now, if I was to use something like barium acetate, that wold be far more soluble in water (about 66 grams per 100 ml vs. 37.5 grams per 100 ml) and this would mean I could add more barium ions to less water, and therefore get more barium hydroxide out of the solution (due to less water, and more barium) So, my reaction would look like: Ba(CH3COO)2 + 2 NaOH --> Ba(OH)2 + 2 NaCH3COO Does anyone know if any of this would be possible?
  18. From what I understand, the tetrahedral arrangement of the phosphorus atoms puts extremely high strain on the bonds - this makes them more reactive.
  19. You're absolutely correct - and it's too bad you're right.... nothing personal, it just sucks that having chemistry as a hobby is so hard -
  20. Right, but as I said, I won't be heating the solution, as that would lead to an explosion.... as far as I am aware, the peroxide won't detonate unless heated...
  21. ... actually, no that is not correct - I work in a school lab with my teacher, who is intimately familiar with it's properties - not to mention I have full gloves, goggle, apron, jump shower and vent hood at my house so even if I wasn't at school, I would still have appropriate equipment Perhaps we should have complete information before we go making snide responses, eh?
  22. Sorry to hear it didn't work - it would be awesome if there was a way to make phosphorus in home labs... Merged post follows: Consecutive posts mergedJust as a quick idea, if there was a polar solvent that could dissolve P2O5 and conducted electricity, you could use electrolysis to isolate the phosphorus, but I'm not sure such a solvent exists.... and water won't work, because it reacts with P2O5 -
  23. That is not correct - there are many chemical suppliers that will sell to individuals -
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