woelen

Most politicians are not interested in the needs for a few chemistry hobbyists, they are only interested in their electroral position. 99.9+% of all people in the USA will think this is a very good thing. It gives them a false sense of safety. They think that terrorists now will have a much harder time making their bombs. The only thing which is reached with this kind of actions is that the people who want to experiment, who grow interest in science, will not be able to find their materials anymore . The terrorist finds the needed materials anyway. Do you really think that a terrorist buys 1 lb of KMnO4 from such a supplier to make his explosive? If he needs KMnO4 then he buys it at the 100kg-scale from some bribed supplier or he arranges something with a corrupt politician. He finds it anyway, cost is not important for such people....

No, budullewraagh, you're not right. Your 80 moles of monoprotic acid will indeed neutralize the 1 mole of OH(-) ions of your CsOH, but after this, the solution definitely will NOT be neutral with a pH equal to 7. Still, the solution will be quite acidic (in fact, the amount of H(+) ions still will be fairly close to 1 mol in this example). On the other hand, if you want to introduce the concept of potentially available H(+), then you could say that there is 80 mol of potentially available H(+). Now, if you add 80 mol of CsOH to this, then again, the solution will not have a pH, equal to 7. In that case, the solution will be quite alkaline and you'll have 1 mol of free hydroxide in solution (disregarding the fact of dilution, but that does not change the story very much). Try to perform the computations with the Ka of the monoprotic acid and you'll see what I mean. Also try something similar in reality with e.g. vinegar and sodium hydroxide.

No, unfortunately you are not right. Things are quite a lot more complicated. What you are stating here only is valid for a strong acid and a strong base, which both are fully ionized in water (e.g. HCl and NaOH). A counterexample: Suppose we have 1 liter of dilute acetic acid (vinegar) at a concentration of 1 mol/l. Such a solution contains much less than 1 mol/l of H(+) ions, somewhat less than 0.005 mol/l. Now, suppose we add 0.005 mol of NaOH to this solution. When all is mixed well, the concentration of H(+) ions still is around 0.005 mol/l, it only has changed around 0.5% and the solution definitely is not neutralized. Computing the amount of NaOH, needed to make this solution have a pH equal to 7 is amazingly complex, but of course it can be done. If you add 1.00 mol of NaOH, then the pH will be well above 7 (somewhere between 11 and 12) and the liquid is quite alkaline. So, 0.005 mol of NaOH does not neutralize, nor does 1.00 mol of NaOH. The same holds for solutions of equal volume. Suppose we have half a liter of 1 M acetic acid and we have half a liter of 0.005 M NaOH (which contain approximately the same amount of H(+) and OH(-) ions, with slight excess of OH(-) ions), and then we mix both liquids, then the resulting liquid of 1 liter of volume still contains H(+) ions in a concentration not far from 0.005 mol/l. If on the other hand we mix 1 M acetic acid and 1 M NaOH at equal volumes, then the resulting liquid will be alkaline with a pH between 11 and 12 (conc. of OH(-) around 0.005 mol/l).

You won't get K-metal at the cathode. Any K-metal formed will react imemdiately with the molten KNO3 around the cathode. Around the cathode you can expect N2 gas to bubble out of the melt and K2O is formed, which remains in the melt. Around the anode, you can expect NO/NO2/O2 to bubble out of the melt. NO3(-) is decomposed to NO2 and O2 and NO2 in turn will decompose to NO and O2 at the high temperatures involved, but I'm not sure to which extent this decomposition will occur.

If you really want to do home chemistry, go for second hand lab equipment. Prices of second hand materials are MUCH lower and for home chemistry it is perfectly suitable. The same is true for chemicals. If you can find an old lab, closing down or disposing of old chemicals, then you're very lucky. Most chemicals are very good, even after tens of years of storage in an old lab. In this way I have obtained many of my chems. Many of them are from old formed GDR-labs for a fraction of the price, which normally would be asked for them.

No, it is not for crystallisation purposes, but you need excess acid in order to dissolve all silver. The oxidizing power of nitrate ion strongly depends on the pH of the solution. Look at the Nernst equation of the oxidizing power of nitrate as function of pH and concentration. Then you'll see that its oxidizing power almost has gone in aqueous solution for any pH well above 0. Only at appreciable concentration and sufficiently low pH, nitrate ion in aqueous solution is a strong oxidizer. So, below a certain concentration, its oxidizing power simply is not enough anymore to dissolve silver metal.

This is NOT an equilibrium reaction. You simply need to add silver metal to excess HNO3. Using the precise stoichiometric amounts does not work, you need quite some excess HNO3. Heating the mix certainly helps, but even then, an excess amount of HNO3 is needed. Fortunately, AgNO3 crystallizes very well and heating of the AgNO3/HNO3 mix certainly gives you nice and dry crystals of AgNO3. If you want really pure (non-acidic) AgNO3, then you need to add the crystals of AgNO3 to a small amount of distilled water and then heat to dryness again. This will drive off virtually all HNO3, which was trapped in the crystals in the first batch. With "heating to dryness" I mean careful heating. First heating, such that a wet slurry appears, with some white solid and then letting all remaining liquid evaporate on a hot room-heater, by letting it stand undisturbed for a day or so. Put a thin paper tissue above the liquid, in order to avoid all kinds of dust entering the liquid while it is drying. A word of warning: AgNO3 really stains your skin and it destroys most organic matter, itself becoming black. Once, you have dry powder/crystals, store it in a small well-stoppered glass bottle.

Well, it's not that pale... Have a look at this picture: http://woelen.scheikunde.net/science/chem/compounds/chlorine.html This is 300 ml of chlorine gas, made by adding dilute HCl to Ca(ClO)2 and leading the gas into this bottle. Quite some more Cl2 was made than this 300 ml, just to be sure that inside the bottle I almost have 100% pure Cl2. This gas, however, was not purified by bubbling it through water and drying it. So, it may contain some HCl and H2O, but these are not visible of course. I never was brave enough to make this bottle full of ClO2 gas, it must be a wonderful thing to see, but I'm too afraid of a possible explosion . Jdurg, did the chlorine gas become more pale, when you cleaned it with water and NaHCO3? That would be a very interesting phenomenon. Indeed, the gas is very pale, when viewed through a tube of 0.5'' width, but if you look at 3 inches of this gas, then it has a nice green color.

It is one of those urban legends that any amount of ClO2 explodes with the smallest amount of daylight. It can explode, when daylight hits ClO2, but this does not mean that it does explode. Besides that, ClO2 has a lower explosion limit of 10% gas by volume. An air/ClO2 mix with less than 10% ClO2 in it cannot explode. Many people confuse an air-mix with a low concentration of ClO2 with pure chlorine. As you can see on my site, ClO2 has a very intense color, while Cl2 only has a weak color. The mix I show on my website definitely can explode, that is why I tell to NOT stopper that test tube. An explosion then results in flowing of the gas out of the test tube (it only is a very small amount), otherwise the glass of the test tube may be shattered around.

I posted the page on setting up a home lab in one of your other threads. In this way you don't get hydrogen and oxygen, but hydrogen and the very poisonous chlorine. Only the anode (positive pole) needs to be graphite, the cathode may be copper wire or any metal piece.

What is wrong with acid spitting in your face? Or even better, with acid spitting in someone else's face?

This is the prototype k3wl-question...

Pyro-things you cannot do with ammonium sulfate. That requires ammonium nitrate, but ammonium nitrate also is not that suitable. It is very hygroscopic. If you really want to do pyro-things, then start with potassium nitrate. This can be purchased as fertilizer (potassium saltpeter, a.k.a. kali salpeter in non-english speaking countries). This fertilizer contains up to 94% of KNO3 and for first basic experiments this is good enough. Wait until spring, in winter time it may be hard to obtain this material and some persons may think it is suspicious if someone buys this stuff in winter time.

Akcapr, read what is posted already in this thread. You did not have NaClO.

Forget about it. Too much impurities. It will be a real pain to extract the ingredients at reasonable purity. If you want pure ammonium nitrate, you'd better buy some fertilizer, based on this compound (together with some chalk) and then dissolve this and let all insoluble matter settle and evaporate the clear liquid to dryness above a warm heater at 60C or so.

Recently I have done quite some experimenting with copper ions and also with iron ions. Now I decided to combine the two metals and a whole bunch of new and special phenomena pops up . Many of you probably know the FeCl3.6H2O in the form of small pea-sized globules, used for etching copper from printed circuit boards. Especially when combined with some HCl it works quite well. If you add some copper to a solution of FeCl3.6H2O in conc. HCl, then the copper dissolves and the solution darkens. Most people think that the following reaction occurs: Cu + 2Fe(3+) --> Cu(2+) + 2Fe(2+) Well, I found that the real reaction is MUCH more complex! What do you expect to happen, when CuCl2 (containing Cu(2+) ion) is added to a solution of FeCl2 in HCl? If you look at the above equation, then you would not expect a reaction at all. In reality the solution turns black or very dark brown. If someone has access to FeCl2 (or FeSO4) and also has access to CuCl2 (or CuSO4), then dissolve some of the iron (II) salt in conc. HCl (appr. 30%) and also dissolve some of the copper (II) salt in conc. HCl. Then add the two solutions to each other. You'll be surprised that the liquid turns almost black. So, there definitely is a reaction. So, apparently Cu(2+) oxidizes Fe(2+) to Fe(3+), itself being reduced to Cu(+). Another experiment: Add a small amount of copper wire to a concentrated solution of FeCl3.6H2O (from an electronics parts store) in conc. HCl. The liquid becomes amazingly dark, almost black, the copper dissolves. Even with a large excess amount of FeCl3 the liquid still becomes black. So even a large excess of FeCl3 does not oxidize all copper to Cu(2+), but part of the copper remains in solution as Cu(+). Finally, add a small amount of CuCl2 (or CuSO4) to a large excess of a solution of FeCl2 (or FeSO4) in conc. HCl. Again, the solution becomes almost black. Conclusion: Excess Fe(3+) ---> Not all copper (I) is oxidized to copper (II) Excess Fe(2+) ---> Not all copper (II) is reduced to copper (I). This seems a contradiction, so there must be a copper (A) species in solution with 1 < A < 2, in other words a fractional oxidation state of copper or a mixed valency complex of copper. This seemingly very simple reaction raises a lot of questions and a lot of interesting things to research. If anybody has ideas or is willing to repeat the experiment and play around with it, you're welcome .

Reactivity is not only a function of electropositivity (or the position in the electrochemical series). Reactivity also depends on things like solubility. A metal like Al is VERY reactive, but still it is stable, because the oxide layer protects the underlying metal. Calcium also is very reactive, but Ca(OH)2 is only sparingly soluble and this makes the reaction between calcium and water somewhat sluggish. It makes the reaction between the water and calcium metal more difficult, because it blocks free access of water to the metal. If NaOH were slightly soluble in water, then Na would not react violently with water. It is the combination of electropositive nature and the perfect solubility of the reaction product, which makes Na (and K and heavier alkali metals) so reactive. The effect of protective layer formation due to insoluble or sparingly soluble reaction products is nicely demonstrated with aluminium, when it is scratched and a small amount of mercury is put on the scratched area. The mercury forms an amalgam and this prevents the formation of an adhesive oxide layer. The Al is oxidized, but the oxide does not stick to the metal and hence the metal is corroded quickly. You can actually see it 'rot away', especially when it is humid.

KFC, your question does not look very smart. Why do you want to make all these materials? I understand that you want to do nice experiments, but if that is what you want, then ask yourself other questions and try to get some knowledge about the subject. So, instead of trying to find out how to make all those chems, try to learn yourself some basic chemistry and buy a good book on the subject. Especially older books can be a really rich source of information about common and particular compounds. Try to obtain some of these. If you have some knowledge, then you'll discover that you want completely different chems than the ones you listed in your post. Making the chems from your list is hard or impossible for the home-chemist without knowledge and equipment. The only feasible ones from household materials are oxygen (from H2O2) and hydrogen (acid + zinc) or electrolysis of water with a suitable electrolyte and suitable electrodes.

Another issue is that kewls are reading this as well. Such persons should not have access to chemicals, because that will cause accidents. I think this is more of an issue than that people want to keep their sources secret. We all benefit from sources, who sell a lot. That makes the companies run better and hence increases the chance that the sources are still there after a few years. But on the other hand, we only want purchases from such sources by trustworthy people, so that it why we should not tell all kinds of addresses in public. So, the intent should not be to keep sources secret, but to share sources in a wise and careful manner. KFC, if you appear to be really trustworthy (depends on your posts and replies in the near future), then people certainly are willing to give you addresses of suppliers. Not through the forum, but through PM. Explain what you want to do with the chems, what kind of experiments you want.

With copper you get copper (II) nitrate. The 3M HNO3 is a sufficiently strong oxidizer to oxidize copper metal to copper (II) ions. The same concentration of H2SO4 or HCl is not capable of doing this. Nitric acid reacts as a stronger oxidizer than the other acids, because the nitrate ion also acts as oxidizer. The nitrate ion is coverted to NO and/or NO2 in this redox reaction. In the other acids only the H(+) ion is the oxidizer: 2H(+) + 2e ---> H2 Iron also dissolves in sulphuric acid and hydrochloric acid. Here the simple H(+) ion is sufficiently strongly oxidizing to dissolve the metal. When iron is dissolved in HNO3 then again also NO and/or NO2 are formed. These form a deep brown complex with iron (II) ions and formation of this complex causes the brown color. This brown complex is [Fe(NO)(H2O)5](2+). Iron is in the +1 oxidation state in this complex, the NO(+) ligand contains N in the +3 oxidation state. So, with iron and nitric acid the reaction is much more complicated than with other acids. In the other acids you simply get H2 and Fe(2+), where oxygen from the air oxidizes Fe(2+) further to Fe(3+).

If you want useful answers, be more specific. Of course from any magnesium compound you can obtain the magnesium, but the effort required to do so may be very high. In general, making magnesium out of its salt, is VERY hard for a simple home-lab. You need very high temperatures and/or a very strong reductor. Using molten MgCl2 (made from MgSO4 by precipitation of Mg(OH)2 with NaOH and then dissolving the Mg(OH)2 in dilute HCl and then evaporating to dryness) probably is the easiest way, but I promise you, this will be next to impossible also at home without the proper equipment. So, practically speaking I would say: NO.

It is soluble in a concentrated NaOH solution. This is due to formation of so-called plumbate (II) anions. Lead (II) is amphoteric. It can react as a base, but also as an acid. Many metal-ions are amphoteric (other examples are Al(3+), Cr(3+), Sb(3+), Be(2+), and even Cu(2+) to a little extent). PbCl2 + 2OH(-) --> Pb(OH)2 + 2Cl(-) Pb(OH)2 + 2OH(-) --> PbO2(2-) + 2H2O The solution has to be fairly strongly alkaline before Pb(OH)2 dissolves. I do not know how much Pb(OH)2 dissolves, but I think a fairly large amount.

Probably the short circuit current has removed a bad contact. It is a fairly well known thing that bad contacts with a large contact resistance become very hot with large currents. With these large currents, the contacts can somewhat melt together and then the resistance drops considerably.

I looked up one of my old books on synthesizing chemicals. Here is a method of making solid hydrated NaClO. The stuff definitely is not what you obtain by distilling bleach. The solid NaClO.5H2O melts at 18 C. A scanned page of the book is included as attachment on this post. This is a scanned page from the book "Handbook of preparative inorganic chemistry", by George Brauer, translated into English by Reed F. Riley. Academic Press, New York, 1963. This is the kind of books from which I get my information and with the info from such books I already have made many interesting compounds. But NaClO.5H2O is beyond my reach .

No, I do not mean a thermite reaction. The OP asks about KNO3:Al mix or KNO3:Mg mix. Many mixes are somewhat hard to ignite and sometimes also have some difficulty to keep them going. If you mix some sulphur with a composition, then it makes ignition much easier. That is why BP has some sulphur in it. A plain KNO3:C mix is quite hard to ignite, with a few percents of sulphur it can be ignited easily with a spark or a simple flame.