woelen

Charcoal from charred wood contains quite a lot of salts (potassium carbonate and possibly others as well), besides carbon. Probably these salts play an important role. I know of a recipe of KNO3, K2CO3 and S, which burns explosively (even when not confined), while a mix with only KNO3 and S does not explode on ignition, when it is not confined. Apparently the K2CO3 plays an important role in this reaction, but I do not fully understand this. K2CO3 is not an oxidizer at all, it contains just spent carbon (CO2, bound to K2O). I think that the K2CO3 in charcoal plays a similar role.

Ryan, I think you misread the question. Acetone and water completely mix with each other in any ratio. If you have a mix of chloroform, acetone and water, then a fairly complicated situation arises from that. Chloroform also is somewhat soluble in water. So, what the precise outcome is, depends on the ratio of the chemicals. But now a practical situation: Suppose you have 10 ml of chloroform, 50 ml of water and 10 ml of acetone and you shake very well. Then two layers are formed, with chloroform at the bottom and water above. The acetone will partly dissolve in the water and partly in the chloroform. Which part dissolves in the water and which in the chloroform is hard to say beforehand, but the amount of acetone will not be negligable in any of the other solvents. The aqueous layer also will contain some chloroform and the chloroform layer also will contain some water. When also an ionic compound (e.g. a salt, or a soluble hydroxide) is added to the water, then the solubility of acetone in the aqueous layer decreases (try adding acetone to a concentrated solution of NaOH, then you'll see that two layers are formed) and more of it dissolves in the chloroform layer.

Sodium permanganate and lithiumpermanganate are not really different from potassium permanganate. The main difference will be the solubility in water and the level of hygroscopic behavior. Sodium permanganate can be obtained commercially, but it is hardly used, due to its hygroscopic properties. The only place, where sodium permanganate is used instead of potassium permanganate is where with the potassium salt the concentration of an aqueous solution cannot be sufficiently high.

What I observe is that all colored substance is absorbed by the charcoal. Some of the aniline may also be absorbed, but certainly not a significant part. I may be because the aniline is not present as the free base, but as its acid salt, in the form of anilinium ions, C6H5NH3(+), while the colored impurities probably are present as large neutral molecules.

I developed a program for balancing chemical equations. A beta-version of this program already was available, but now I present a new version, which allows the use of preserved groups (useful for organic chemistry) and the use of isotopes. The program can be used for balancing any chemical equation, e.g. it nicely balances equations like the following: [ce]..K2Cr2O7 + ..KCl + ..H2SO4 -> ..KHSO4 + ..Cr2O2Cl + ..H2O[/ce] The program determines the coefficients, needed to make it correct: [ce]K2Cr2O7 + 4KCl + 6H2SO4 -> 6KHSO4 + 2Cr2O2Cl + 3H2O[/ce] The program can also do computations from moles to grams and vice versa and it computes weight ratios, needed for certain reactions. It also detects when a chemical reaction is non-stoichiometric (a nice example is the reaction between copper and nitric acid, where both NO and NO2 are formed) and the solution cannot be determined unambiguously. The program also allows the use of ions in the equations, so it can be used nicely for balancing equations, in which dissolved ionic compounds are involved, e.g. ions like Cl(-), Cr2O7(2-), MnO4(-), etc. If you download the program, please also read the tutorial, before asking questions to me about the program. Any feedback on the program, functionality, user interface, algorithms, is welcome. The program is available at this page: http://woelen.scheikunde.net/science/chem/chemeq/index.html It is very easy to use, no installation needed, just click the icon and start playing....

This evening I had a small explosion, while demonstrating this reaction to some friends. I performed the reaction with a drop of after shave. I did this with the idea to use something they all know as a common substance. As soon as the drop of after shave fell on the Mn2O7 it exploded with a loud and high pitch and a violent bright purple flash and a lot of very small droplets of Mn2O7/H2SO4 were sprayed around, unfortunately also some on my hand and clothes . Lateron I repeated the same experiment with acetone, that again gives more like a flame as shown in the small animation. So, the nature of the reductor has a strong influence on the outcome. Please be careful! It is surprising to see that 50 mg can do so much harm and creates such a loud and scary explosion (one of the persons was not amused at all ).

I use activated charcoal for making aniline colorless, before I use that for experiments. Aniline (C6H5NH2) has a strong tendency to become brown on storage, due to formation of all kinds of oxidation products. When the aniline is added to some dilute sulphuric acid and dissolved, then the resulting solution of anilinium sulfate (C6H5NH2)2.H2SO4 is brown/red. With some activated charcoal added to this, it becomes colorless. What effect is responsible for this? Can the brown material also be released again by the activated charcoal?

Have a look at this link: http://science.csustan.edu/stkrm/Recipes/Recipes-Blue.htm Indeed, it is a nice demo.

Jdurg I did the experiment you suggested. I have some reagent grade powdered graphite and it does NOT react. The charcoal I used for the first experiment probably is much more reactive, due to unterminated carbon atoms, while the graphite has a much better defined structure, with all carbon atoms (except at the edges of the 2D sheets) being fully incorporated into a stable structure. I first need to rob a jewellery shop before I can do the last experiment . Lately I just was running out of diamonds .

The answer of akcapr is correct. There only is very little color in the solid soap. Due to scattering of light in the foam, the color becomes less visible. The finer the foam, the less of the color can be observed. A nice experiment which demonstrates this scattering effect is to take a small piece of colored soap (or even better, take a small piece of a strongly colored hard and brittle sweet). Crunch the piece into a fine powder. The finer the powder, the lighter the color. This effect is the same as with the foam. Because there only is such a small amount of dye in the soap, any coloration of your skin will be unnoticed.

As Ryan already stated, please read the page which you found yourself. It gives all information you need. And no, you cannot replace the zinc with sulphur. Then it will not work anymore. Still, you can make it burn, but it needs ignition then with quite some heat. KFC, please do not post all kinds of random question on this forum as you do right now and as you did before your temporary ban. Please do some research yourself. People are willing to help you with your questions, but you do not show a very intelligent attitude over here. Shooting random questions without showing any progress in your postings does not look very good. Try to be a person, who goes beyond the k3wl-level. Next time better?

Calcium salts are very easy to get your hands on. E.g. for drying purposes you can buy CaCl2. Purity is not even an issue. The CaCl2 usually also contains some NaCl and MgCl2, but none of these interferes with the test. Another alternative is calcium carbonate. Very easy to obtain in pure form. Simply dissolve some of this in dilute HCl and filter the solution to get rid of solid particles. Use a small excess of CaCO3 in order to make the solution as little acidic as possible. With this method of precipitatation it is very important that all solutions you use are perfectly clear. Any solid particles should be filtered out, before you make the precipitate of CaF2. Having solid particles in your liquids affects the outcome. They will also be centrifuged together with the precipitate.

Akcapr, did you ever hear of the pot and the kettle ???? Look at your own posts almost one year ago. But keeping things serious. Please don't flame other members and don't put yourself on a higher level than other members. ----------------------------------------------------------- KNO3 is an ionic solid and when it dissolves in water, it splits into K(+) ions and NO3(-) ions. These ions are hydrated. In fact, KNO3 is actually two different compounds (entities) in a single crystal lattice. This is very nicely demonstrated when two different salts are dissolved in water. E.g. dissolve 1 mol of NaCl and 1 mol of KNO3 in a liter of water. Then you obtain Na(+) ions, K(+) ions, Cl(-) ions and NO3(-) ions in solution. If, on the other hand you dissolve 1 mol of KCl and 1 mol of NaNO3 in a liter of water, then you obtain exactly the same solution. Noone can ever distinguish between the two solutions, because they really are the same!

Of course this problem can be solved theoretically, but you have to make an important assumption. The solid KCl must be completely compact (no air between the granules). Otherwise a statement about ml's of this solid is meaningless. Usually, with solids we talk about weight and your question then can be restated as follows: "I was wondering if there is some kind of equation out there to solve a problem like this. How many milliliters of a 25% solution of potassium chloride must be added to ... grams of pure potassium chloride to obtain a 35% solution." You need to know the density of a 25% solution of potassium chloride. You can measure this, or use some published table. With that info the problem is simple. Suppose you have X grams of KCl, then you proceed as follows: Suppose you have Y grams of a 25% solution of KCl and you add X grams to this and all of X dissolves (assuming that KCl is so soluble that a 35% solution can be made), then you have a liquid with (0.25*Y + X) grams of solution. The total weight of this is Y + X grams. So, you need to solve the following equation: 0.25*Y + X = 0.35*(Y+X) Solving this is piece of cake, for a given Y.

The concept of stability, introduced here, does not make any sense. There are other elements, which are much more stable than carbon (with stability in the sense of redfox), such as gold, platinum. What is special about carbon is that it can form tremendously large and complex molecules of great diversity. No other elements even shows a pale shadow of carbon's great versatility in forming different compounds. With other elements, we already are quite surprised if we find clusters of the same element with bonds between atoms of the same element of more than a few atoms. Such clusters exist, such as Pb-clusters with around 10 atoms, where Pb-Pb bonds exist. The same is true for Re, B, and to a lesser extent Si. But carbon can form almost any type of structure with hundreds or thousands of carbon-carbon bonds in it. This makes carbon unique and this property is exploited by all life-forms we know.

Another method may be to precipitate all fluoride with a calcium salt. This forms the very insoluble CaF2. With a centrifuge the precipitate can be collected and rinsed with distilled water. Again centrifuging and rinsing with distilled water and then a final centrifuge. Finally, the precipitate has to be heated in an oven at 200C for an hour or so. This removes all traces of water. The precipitate can be weighed. This method of course requires a centrifuge and a very accurate scale.

I did the experiment with carbon, sulphur and glycerol. With carbon, a small explosion occurred. This is not very spectacular, when viewed in an AVI file, but the experiment is quite nice. As soon as the piece of carbon falls in the liquid, there is a fairly loud crackling noise and a bright flash. An AVI file of this is presented here anyway, but the result is not as nice as with the acetone or alcohol. http://woelen.scheikunde.net/science/chem/exps/mn2o7/mn2o7-c.avi I also did the reaction with sulphur. Remarkably, with sulphur there was no visible reaction. The material did not set off, no fire could be observed. Finally, I did a test with glycerol. The result with that also is quite nice. An AVI file of this is available here: http://woelen.scheikunde.net/science/chem/exps/mn2o7/mn2o7-glycerol.avi

YT, your ideas are interesting. I'll repeat the experiment with solid carbon on Mn2O7. I have some activated charcoal lying around and I could try putting a few granules of this in Mn2O7. The same holds for sulphur. I do not expect that passing a flammable gas over this will ignite it, the concentration of reductor in a gas is much lower than in a solid or a liquid. @Ryan: Why do you think that your gloves will not react with this material? What are they made of? It would indeed be interesting to see whether this material reacts with the rubber of the gloves.

Yes, using the pre-mixed in the way as YT suggests is perfectly OK with me. I sometimes do that kind of experiments with my children. They really like it. Last time I did this, I made a mix of KBrO3 and S and I gave them a magnifying glass, mounted on a long stick and they were allowed to light the KBrO3/S mix by focussing the light of the sun on the mix. They loved it!

Caffeine.HCl is a chemical compound, which inherently contains HCl. It is a hydrochloride salt of caffeine. Many compounds form such salts, e.g. hydrazine, N2H4.2HCl, cocaine.HCl. When they are heated, then the HCl is released again and that is what you smell. The reason that many compounds are made into an acid salt is that these are much more resistant against aerial oxidation and they are well-crystallized solid compounds. This makes storage and transport easier.

I do not agree. I have no objection against that particular reaction between S and KNO3. My objection is the way it is performed. I keep my opinion that it is plain stupid to mix the two, put the two in a test tube and then heat the mix. I have given a few suggestions on how it can be done more safely, so why would you want to do it in that unsafe way? The reaction looks the same (as spectacular) when sulphur is added afterwards, while the risk of chemicals being sprayed around is substantially lower.

Why do you want a real Bunsen burner? A well-adjusted propane torch is as safe and also perfectly does the job. I unfortunately cannot help you with an address where you can buy a real Bunsen burner, but if you purchase one, be careful to select one for the proper gas. A burner, made for buring natural gas is not suitable for use on propane bottled gas. The same is true for a burner, designed for propane. This does not work nice on natural gas. The flames become very ragged and uneven.

It IS dangerous. What akcapr describes is adding both compounds to a single test tube and then melting. In this way you get a mix of chemicals and when the KNO3 starts reacting with S somewhere in the lower part of the test tube, a lot of gas is produced (mostly N2, SO2), and this gas may swirl the material above it out of the test tube. This simply is a stupid experiment. If you want a (still quite dangerous) but relatively safe experiment, then put some KNO3 in a test tube and heat this, until it is molten. Then add some sulphur. The effect is the same, but MUCH more safe. Now the reaction occurs in the top of the test tube and the chance of molten chemicals being sprayed around is much smaller. Still, however, this experiment also must be performed with great care. The test tube may crack due to the heat and still some hot material may be swirled away from the test tube.

I also use a propane/butane burner. It is a very cheap thing and it perfectly suits my needs. It also has a wide bottom, so that it cannot easily be bumped. Have a look on this page for a picture and description: http://woelen.scheikunde.net/science/chem/misc/equipment.html The burner, including three gas tanks, only cost me around $20. The intensity of the flame can be adjusted precisely and usually, not the maximum power is used, but a much lower setting. With that setting, I only need one gas tank per year or so, while I do quite a lot of experiments, which require heating of a test tube.

This is a very dangerous thing to do!!! Never melt a mix of KNO3 and a reductor in a test tube. The reaction may become so violent that the contents is swirled away, out of the test tube, spraying around molten chemicals! You can do the experiment with KNO3 and S by mixing it carefully and lightling the mix in the open on a piece of stone or metal. But never do this in a test tube! @akcapr: Next time, please be more careful with what you suggest. Nobody is waiting for accidents with burning chemicals sprayed around!