woelen

MEKP also is quite unstable, but less unstable than most other organic peroxides. The question, why compound A is stable, while compound B with similar features is unstable (or less stable) cannot be answered easily. With advanced mathematical models of molecules, based on quantum mechanics, one can predict stability by means of a 'simulation'. But for most compounds we know, it is just a matter of experimenting, because the computations require MASSIVE amounts of CPU resources and certainly are beyond what the home-chemist can do with his PC . Sometimes, by means of reasoning you can make sensible predictions about a compound's stability. Indications for low stability are: 1) the presence of extremely high or extremely low oxidation states for a certain element (e.g. manganese (VII) compounds, sodium electride). 2) large mismatches in atomic size (e.g. CI4, with small C atom and 4 large I atoms) with multiple large atoms around a small atom 3) the presence of peroxo and superoxo structures in a compound 4) the presence of oxidizing groups in covalent compounds (best known examples of these are nitrate-esters, chlorate-esters, permanganate-esters and perchlorate-esters). 5) the presence of steric hindrance, which causes strain in molecules

When mixed, these give NH3, H2O and BaCl2. I would expect this to be an exothermic reaction, although not a really strong exothermic reaction. Endothermic reactions are not dangerous at all (at least not in terms of violence, heat, ignition, explosion etc.).

Not all organic peroxides are extremely unstable. An example of one, which has a commercial application in the manufacture of resins is MEKP (methyl ethyl ketone peroxide). But, still most of the organic peroxide are very unstable and really dangerous. This has to do with the peroxide bond. A single O-O bond is very weak. It is easily broken. The resulting radicals are VERY reactive and in this way an avalanche effect can be created very easily. Not only organic peroxides are notoriously unstable. Just think of plain H2O2. At higher concentrations, this also is quite a dangerous chemical. Even the commercially available 30% grade material already has to be handled with due care! Many chemical reactions with H2O2 must be carried out at low temperatures.

The element copper remains a very special element to me. It keeps on giving strange results and I want to post one over here again. It is a really simple experiment, just add two chems to each other. In jsatan's words, I saw a "little thing", which I cannot neglect. The "little thing" here is that the copper (II) chloride turns very dark brown, almost black in the concentrated acid. When water is added, then it dissolves again and the results become predictable again. http://woelen.scheikunde.net/science/chem/riddles/copper+h2so4/index.html If anyone over here has an explanation for the observation, then I would be pleased to know.

Never ever pour water on a running thermite reaction mix!!! This results in instant explosion, with molten metal sprayed around. I once read a story about someone putting a zinc coin on the mix. That also leads to explosion. The heat, produced in the reaction causes the zinc to vaporize at once and the instant expansion of solid zinc to gaseous zinc causes an explosion!

Yes, I understood that water can burn in fluorine! That would be really cool to see.

HNO3 can even be 100%. This cannot be made by distillation of dilute HNO3, due to formation of an azeotropic mix of 68% HNO3/32% water, but with other methods, 100% HNO3 can be made. H2SO4 also can be 100%, but commercial H2SO4 has at most 97 .. 98% acid. This also is due to the formation of an azeotrope with water. Also H3PO4, CH3COOH and HCOOH can be at concentrations of 100%, although the commercial acids usually are at 85, 80 and 85% respectively. The limit of concentration of HCl is that pure HCl is a gas and the solubility of this gas is limited. Max concentration at standard air pressure is approximately 38%. It is as with carbonated softdrinks. At higher pressure more gas can be dissolved than at lower pressure. If the concentration of HCl would be higher than 38%, then the bottle would be pressurized, just as with carbonated softdrinks and as soon as such a bottle would be opened, gaseous HCl would bubble out of the solution. I have a bottle of 37% HCl and on warm summer days, there already is quite some over-pressure in this bottle (which is a real pain, because everything near this bottle is corroded ). The acid HF can be 100% and at that concentration it is a liquid, which boils at just under 20 C at standard air pressure. Pure HF indeed has a concentration over 35M, the molarity is the density of this, divided by the molecular weight of this. For water, the molarity even is 55M.

Fortunately we all are human including RyanJ

To my opinion Brainiac is sheer crap. The only thing the program does is grow k3wls. It may be fun to watch, but the scientific contents is 0.0, or even less than 0.0 (i.e. sheer nonsense). Btw, NaCO3 would be a rather interesting chemical which I certainly would like to do some research on . What you probably meant is Na2CO3 .

If you have more of these "little things", I would invite you to post them if you don't have a (trivial) explanation for them. In this way, I have come up with quite a lot of riddles already (I have some on my website in the riddles section). I also am very picky on those little details. Such little details sometimes lead to very surprising insights or discoveries.

Forget about making CS2 at home. This is exceedingly difficult. Have a look on http://www.sciencemadness.org and search for this subject. It is in a recent thread. If you want to dissolve sulphur, then you can use hot toluene. Almost pure toluene can be obtained at hardware stores as thinner. It's a nice and MUCH safer alternative for CS2. Sulphur dissolves in warm thinner quite well, in cold thinner it dissolves in smaller amounts. I did the experiment and put it on my website. http://woelen.scheikunde.net/science/chem/exps/S+toluene/index.html If you want to repeat this experiment, please heat the toluene with boiling water and not with an open flame. Toluene is quite flammable and it is moderately toxic.

This question is not easy to answer at all. How atoms are arranged around each other depends on the number of electrons, the level of hybridization and the presence of lone pairs. For methane, CH4, the four H-atoms are at the vertices of a regular tetrahedron (think of a pyramid, but with a triangular base). In methane, all orbitals of the C-atom are [math]sp^3[/math] hybridized, making essentially four equal bonds. That is precisely what you see in the tetrahedral CH4 molecule, all H-atoms are equivalent. If you really want to understand the subject, then I would like to suggest you to read something about orbitals, molecular orbitals, hybridization. As a first start, some basic quantum mechanics is needed.

Indeed, cyanide is not cumulative and it is metabolized fairly quickly. Either you have a lethal dose and it kills you quickly, or you recover from it and all traces of cyanide disappear from the body within hours.

So, there must be at least one parallel universe without YT :D not to count all the other could-be-events

Well, in fact RyanJ, if you experiment a little more, then many of these things you probably will make yourself. E.g. bromine and chlorine I make quite often, just for experimental reasons and I'm actually quite comfortable with these. I know the dangers (and yes, they are dangerous), but this does not need one to refrain from using/making them. But as I stated before, KNOW WHAT YOU ARE DOING. In chemistry it simply is true that many fascinating compounds are quite dangerous. An important criterion for me not doing an experiment is unpredictability (such as with AP and NI3) and lack of information. If I can't find sufficient information about a compound and I've heard/read that it is very dangerous then I don't make it. Well, the following is not good at all and I've been exposed to it far too much unfortunately (I even ingested quite some amounts ): alpha-d-glucopyranosyl-(1->2)-beta-d-fructofuranoside A nasty longterm effect of repeated overexposure to this stuff is that it can make you really fat. So, if you use this, be careful. Unfortunately it is where I live. I have been looking for sources of uranyl salts, but where I live, they cannot be purchased, nowhere, never. I've not the courage to import it from the USA, I'm quite sure that the package will not pass the customs and that I'll get BIG trouble with it. Antimony compounds are fairly toxic, but not very special. At ceramics and pottery suppliers you can buy antimony oxide, which you can dissolve in HCl to make an acidic solution of antimony chloride. Elemental antimony can also be purchased easily on eBay for a reasonable price at amazing purity. I personally have done quite some experiments with antimony, the most beautiful compound I made being red antimony sulfide. Yes, it is. In fact, there are two forms, a dimer and a trimer. The dimer is the most crappy stuff there is. Totally totally useless because of its instability. The trimer is the next most crappy stuff (together with NI3 ).

For me, the most dangerous chems, which I've used in experiments are: 1) Hot anhydrous mix of H2SO4, HF and KMnO4 (see experiment on website). This stuff is insanely corrosive and also very toxic. 2) Hot mix of NaCN and aqua regia. This bubbles nicely, like carbonated softdrinks, the main difference being that the bubbles are a mix of NOx and HCN instead of CO2 . As you see on my list, no explosives like NI3 or AP. I'm not particularly after very toxic, corrosive or otherwise extremely dangerous chemicals, but sometimes coincidently a very interesting experiment also involves very dangerous chemicals. I, however, ALWAYS perform this kind of experiments after doing kind of risk assessment, based on the properties of the chems, but also on the risks of spills, breakage of glass, etc. Sometimes the risk assessment takes many hours on Internet and in books, while the experiment itself only takes minutes. I think that is the great difference between k3wls and serious home chemists. K3wls first act, then think (or repent ), while a real citizen chemist first thinks once, twice or even trice and then acts (or not).

This is VERY good man! Now, you can purify your benzoic acid very easily. Make benzoic acid from sodium bezoate and any acid. The precipitate still contains some sodium benzoate and other ionic stuff. Now you dissolve your (somewhat impure) acid in the ethanol. Any insoluble matter (including remains of sodium benzoate) do not dissolve. Let insoluble matter settle at the bottle and decant the clear liquid. Next, let ethanol evaporate and the benzoic acid remains, free of sodium! Now, you can make copper benzoate, absolutely free of sodium as follows: Prepare a solution of potassium benzoate by dissolving as much as possible of benzoic acid in a solution of KOH or K2CO3. Next, add a solution of copper sulfate to this solution. Now you get a precipitate of coper benzoate with some co-precipitated potassium and sulfate ions. These ions do no harm, because they do not spoil the flame color. Filter the precipitate, rinse a little and let dry overnight. You see? You have a perfectly doable method of preparing copper benzoate, which is absolutely free of sodium ions and you exploit the fact that benzoic acid is soluble in ethanol, while sodium benzoate is not.

Copper benzoate is not soluble, but you'll find it extremely difficult to get it totally free of sodium ions. Also if you make Cu(OH)2, it is extremely difficult to get it in the pure state. A general rule: If a precipitate is formed of a flocculent nature, such as Cu(OH)2, then you'll suffer from co-precipitation. Together with the Cu(2+) ions, some Na(+) will be encapsulated in the precipitate, which cannot be washed out. Of course, the same holds for the sulfate ions. If you precipitate a crystalline solid, then the problem of co-precipitation is much less severe. So, if you want to make copper benzoate for making whistle mix, which has another flame-color than orange, try to make some potassium benzoate and precipitate that with copper sulfate. In that case you get copper benzoate with some potassium and sulfate ions co-precipitated, but these do not spoil the color of the flame. I understand that you made your benzoic acid from sodium benzoate. You can get rid of sodium in your benzoic acid, by dissolving it in an organic solvent, decanting the solution and discarding any solid which does not dissolve (this contains the ionic remains of sodium benzoate) and evaporating the solvent. I'm not sure what is a good volatile solvent for benzoic acid. You could try it with acetone. Even if trace amounts of sodium are left in your mix, it spoils the flame color.

Try mixing strontiumcarbonate and benzoic acid with a small amount of water. Strontium carbonate can be purchased at ceramics and pottery suppliers. You might be able to make strontium benzoate with this and use that as fuel. I'm not really sure whether this method allows you to make strontium benzoate, but you could give it a try. Strontium carbonate is dirt cheap and easy to obtain. Beware, that in commercial strontium carbonate there may be up to 5% of strontium sulfide. So, perform the mixing and preparation of strontium benzoate outside!

All depends on quantities. If you make 20 mg of the stuff and that would explode in your hand, then you would feel quite a bang, but no lost fingers. If a whole gram explodes in your hand, then your entire hand may be torn apart. But, it IS very sensitive and powerful. Btw, I also strongly advice against making AP. This also is nasty stuff and you would not be the first person who becomes seriously injured by that crap. Last summer a Dutch boy had blown of a few fingers and he'll have to do the rest of his life with less than 10 digits. Even unscrewing the cap of a container, in which AP is stored can make the whole bunch detonate. To my opinion, people fiddling around with AP or NI3 potentially ruin their own lifes, but they certainly ruin the possibilities for other home chemists, due to liability issues with suppliers. What is this that all these people want to fiddle around with crap like NI3 or AP? This is not fun at all and certainly not worth the risks.

Can you explain your reasoning to me by a worked out example for 1 M H2SO4? I do not understand what you mean. What do you mean with "molar percentage"? I just cannot get any other percentage than 10% by weight for 1 M H2SO4. If I were to look for the percentage of the number of moles of molecules, then I would find approximately 2% (pure water is 55M, we have around 10% less, so 50M and we have 1 M acid in it), but this is just arithmetic and has no real useful meaning.

No, 1 M H2SO4 is almost 10% by weight. You forget to take into account that conc. H2SO4 has a density over 1.8 gram per ml, while 1M H2SO4 is close to 1 gram per ml.

For the transition metals you can say that the metals of the first row are much better reductors than the corresponding metals of the second and third rows. This is true for all of these metals (i.e. Sc to Zn and all metals in between, including Fe).