woelen

You will have a hard time finding a supplier for white P and potassium metal. If, however, you find some, please do not simply buy these things. White P is VERY toxic (50 mg may kill you, even 5 times as toxic as cyanide). It also ignites when the whether is somewhat warm. I have a small piece of white P for my element collection and while I was transferring it from one container to another it immediately started smoking, as soon as I took it out of its container. This was really scaring ! Potassium is another dangerous beast, which will certainly bite you badly, when you do not know how to handle it correctly. It is MUCH more reactive than sodium metal and reacts almost explosively with water. It also easily ignites in air, making an almost impossible to quench fire. Besides that, it can form peroxides on storage, which are insidiously explosive.

Dissolving any form of carbon is extremely difficult. There are some compounds, in which it dissolves, but then it reacts chemically. I do not know any solvent, which dissolves carbon, without changing it chemically. It is also easy to understand, why carbon cannot be dissolved without changing it. Carbon does not form small molecules, but it forms macroscopic molecules (sheets in graphite, irregular structures in charcoal, 3 regular 3D structure in diamond). Dissolving would mean that bonds have to be broken. A similar thing is true for red phosphorous and vitreous (black) selenium. These also form large macro-molecules and cannot be dissolved. A compound like sulphur has S8 molecules, white phosphorous has P4 molecules and red selenium has Se8 molecules. These all can be dissolved in certain solvents (CS2, toluene).

Formaldehyde is a gas at room temperature. The use of pressurized containers is not an option. Formaldehyde can polymerize easily and such a container would be clogged completely within a few months. Fortunately, this gas dissolves in water VERY well and that is why a solution in water is sold. Also this aqueous solution suffers from polymerization. If you have a bottle of formalin (37 - 40% CH2O), then you'll certainly see a white layer of paraformaldehyde at the bottom of the bottle.

KFC, with careful heating, you can concentrate a solution of H2O2 to a higher concentration. The loss, however, will be considerable. If you want to make higher concentration H2O2, then use a freezer. Freeze some of the 3% solution and let appr. 50% of the liquid freeze. The ice can be thrown away, it is mainly water. The remaining liquid contains almost 6% of H2O2. Repeating this procedure with that 6% liquid will bring you to over 10% of H2O2. In this way, it is said that you could even reach 30%, but I've never tried that. Reaching 10%, however, can be done quite well, with acceptable losses (due to presence of some H2O2 in the ice).

It probably is used in the form of gallium arsenide. GaAs is a semiconductor, which allows faster circuits to be built than silicon. It is, however, much more expensive and it is harder to obtain a large component density. So, it only is used in places where speed is an absolute premium and where cost is a less important criterion.

The transformer may also cause troubles. I have made circuits for driving motors. The inductive peaks, when the motors were switched on could be so powerful that semiconductor devices are blown out. The voltage across an inductor is proportional to the time derivative of the current through the inductor. If the circuit is switched on and a large current is running through it at once, then you could obtain very high peak voltages.

In a pool the typical smell of chlorine is due to a mix of elemental chlorine, hypochlorous acid (these two are desired and are disinfecting), and chloramine, NH2Cl. The irritation, however, mainly is caused by the chloramine. The latter compound is formed from urea (and derived from that, ammonia) and hypochlorites in the pool-water. The chloramine also is not useful anymore as disinfecting agent. It only irritates. In a pool there are two kinds of chlorine: free chlorine and total chlorine. Free chlorine is the Cl2 (and HOCl) in the pool and total chlorine is free chlorine plus the chlorine, bound in NH2Cl. If a pool has a very strong and irritating "chlorine" odour, then surprisingly, a shock-treatment with a lot of hypochlorite makes the pool less irritating and less odourous. With the excess hypochlorite, the chloramine is destroyed (giving mainly N2, water and chlorine) and the irritating stuff is gone.

No, we do not do that. Everyone gets a chance to improve, and in fact, I see improvement already in newer posts . Let's see how things are developing.

Also a word of warning with gallium. It wets everything like hell, when it is molten. It is not like mercury, which forms little globules, which do not stick to anything. Gallium is more like water. Your skin will become wetted by it. It forms a kind of silvery layer, which is VERY hard to remove. And trust me, even 1 gram of gallium is sufficient to make you look like a robot or alien from one of those cheap and ugly SF-movies from the seventies. I certainly would not recommend playing around with gallium like this. You may regret it. One positive thing: gallium is not very toxic, so, even if you look like an alien, you do not have to worry about health effects. IIRC the density of gallium is somewhere around 6 to 7 grams per ml, so, an amount of 5 grams will not be very much. You'll have a piece at the size of a small marble.

Fluorates do not exist. Fluorine is an element, which only exists in oxidation state 0 (as element) or in oxidation state -1 in ALL of its compounds. There is "hypofluorous" acid, HOF, but this in reality contains fluorine in oxidation state -1, oxygen in oxidation state 0 and hydrogen in oxidation state +1 (as it usually is). So, HOF is totally different from the other hypohalogenites, such as the common bleach, which on very slight acidification gives HOCl. Higher fluorates simply do not exist. So, when NaF is electrolysed, then indeed, as Jdurg already stated, O2 and HF are formed at the anode and OH(-) and H2 are formed at the cathode: anode: [ce]2H2O + 4F^{-} - 4e -> 4HF + O2[/ce] cathode: [ce]4H2O + 4e -> 4OH^{-} + 2H2[/ce] Net reaction when the solution is mixed: decomposition of water.

Milk contains some grease, so this indeed is non-polar, but with beer that effect is not present. Beer is an aqueous solution and the capsacain does not dissolve in beer. Indeed, drinking beer is not nearly as effective as drinking milk, when the 'hot' taste has to be suppressed.

I would like to say [ce]C_nH_{2m}O_{m}[/ce], with m ≤ n. Polysaccharides can have a lower m value, but mathematically, m and and n are different numbers, hence the two different values.

Just to make things clear. I have absolutely not the idea that KFC (are you still there?) has any bad intents. He seems interested in the matter, but always in a somewhat kewlish direction. No, not dangerous in the sense of evil intents, criminal behaviour or terrorism, not at all. But somewhat irritating (at least to me), because there seems to be very little progress in what is learnt, at least when we look at the posts on SFN. I would like to invite KFC to read more books on basic chemistry, try to do a little more research on the subjects he posts and to formulate a first thought on the direction he is thinking off. That makes the posts look a lot better.

You only use an N/10 solution? I was speaking about 10% by weight. N/10 is VERY dilute (only appr. 0.35%). With a 10% HCl solution and a current of 0.75A you should have a nice green solution within an hour or so. HCl can be used to make carbon dioxide in large quantities with baking soda. With chlorine bleach you can make green chlorine gas, but beware, the latter experiment MUST be done outside on a windy day, with wind from behind. Chlorine gas is very toxic. The experiment, however, is very nice and chlorine gas can be used for many other interesting experiments (such as burning iron wool, or fine copper wire with yellow/brown smoke).

No, the copper already present makes the reaction much faster. Copper plates dissolve much eaasier in a solution which contains already some copper. Hence, the copper is self-catalyzing. In fact, the solution will become better and better when more copper is in there. You only need to replenish it sometimes by adding a small amount of conc. HCl. KClO3 is not needed anymore. It is important though that the liquid has good air-contact. You get your oxidizer for free, being oxygen from the air. How does this work? Initially you only have HCl in solution. That does not dissolve copper. You add a small amount of oxidizer (in your case KClO3, but H2O2 or even a few ml of chlorine bleach also are OK). This small amount of chlorine, formed from the oxidizer and the HCl is capable of dissolving copper, albeit with difficulty: [ce]Cu + Cl2 + 2Cl^- -> CuCl4^{2-}[/ce] The complex ion, formed in this way, however, is a very good oxidizer for copper. Here appears a very specific piece of copper chemistry. Copper is capable of forming special complexes with multiple oxidation states in one species and this complex breaks down into two complex ions with copper in the +1 oxidation state: [ce]Cu + CuCl4^{2-} -> Cu.CuCl4^{2-} -> 2CuCl2^{-}[/ce] The complex ion CuCl2(-) very easily is oxidized. Even oxygen from the air can easily oxidize this in the presence of acid: [ce]4CuCl2^{-} + O2 + 4H^{+} + 8Cl^{-} --> 4CuCl4^{2-} + 2H2O[/ce] The CuCl4(2-) hence is recycled again and effectively, oxygen is used as oxidizer, and some acid is needed for the oxidizer being effective. It almost seems too good to be true, but still, it is. You only need HCl and oxygen from air to etch copper metal. You only need a small amount of oxidizer (or even better: a copper salt) to get things started. If your liquid turns very dark, then you need to add some acid and you need to let it stand overnight with air contact and then the liquid is replenished again. This very specific reaction only works in the combination HCl/copper. It also works with HBr/copper, but not with other acids, like H2SO4, HNO3 and also not with other metals.

YT, try some of this green liquid with your Al-foil experiment from that other thread on electrolysis. You'll love the result. The small amount of KClO3 in it is irrelevant for the Al-experiment, so using this green liquid perfectly works.

Oh yes, HCl and KClO3 will make a great etching combo, but also a somewhat dangerous one. When the acid is concentrated, then ClO2 is formed and also Cl2 is formed. These fumes are not nice. A safer alternative is to use HCl, with some H2O2 added. This also produces bad fumes, but it cannot make ClO2.

If you want to make sulphur crystals, then you need another solvent. H2SO4 is a very bad solvent for crystallization purposes . A good solvent is toluene at somewhat elevated temperatures. It allows you to make nice sulphur crystals. Another solvent is carbon disulphide, but that is very toxic, extremely flammable (water of 95 degrees Centigrade can ignite CS2) and hard to get. If you want to isolate sulphur from match heads (assuming that some is in them), then water would be the best. Salts in the match head dissolve, sulphur remains behind. But I strongly doubt that your match head contain sulphur. What color do these match heads have? Only if they are light yellow, I can believe they contains sulphur. If they are red or brown, then they almost certainly do not contain sulphur, but some sulfide, combined with an oxidizer and some colorant.

Electrolysis is quite a slow process, if you wait long enough, then the liquid will turn green (or light green/blue). The white material most likely is CuCl, formed at the anode. This is white. It dissolves in the HCl and will be oxidized quickly by oxygen from the air to a copper (II) compound. Once your solution is greenish, you'll have lots of H2 almost at once from your Al-foil. The black coating at the cathode must be copper metal, in the form of a spongy mass. In that way, it looks very dark. Try to avoid as much as possible that the white material from the anode reaches the cathode. What you can do is wrap a paper tissue very loosely around the cathode, such that there is no free flow of liquid around the cathode, but still with sufficient room around it, to let all bubbles escape. Your outcome surprises me, I've done electrolysis of HCl quite some times myself with copper electrodes and it worked like a charm. I must admit, however, that I used concentrated HCl, which I diluted after the electrolysis. YT, do you have some CuSO4 or CuCl2? If you have that, add a pinch of this to some dilute HCl (10-15% is OK), dissolve this, such that you get a green solution and then add the Al-foil. The result is really stunning! If you do the same without the copper salt, then the Al only reacts after a long induction time. Even stronger, this experiment even works with plain NaCl instead of HCl, if some copper salt is dissolved in the NaCl as well! The Al dissolves in the solution of NaCl with a lot of hissing, evolution of a lot of heat and formation of a lot of hydrogen gas! I'm not kidding, this really is true! http://woelen.scheikunde.net/science/chem/exps/cu+al/index.html

Sulphur does not dissolve in sulphuric acid. Only with strong heating you may manage to dissolve some of this, but this is not something for the unexperienced hobbyist. A match head usually contains an oxidizer as well (this might be KClO3), some filler and some antimony sulphide. The KClO3 decomposes when brought in contact with H2SO4 and the decomposition products may explode or ignite on contact with the reductor in the match head. I think it is quite dangerous to do this experiment. There is a risk of concentrated H2SO4 being sprayed around. When you use dilute H2SO4 (such as car battery acid), then nothing will happen at all. Part of the match head will dissolve, the remaining part will form a slurry at the bottom. ------------------------------------------------------------------- Concentrated H2SO4 is capable of oxidizing copper metal, but it has to be warm. Not because of its acidity, but because of the oxidizing properties of the sulfate under these conditions. The acid then does not form hydrogen gas, but sulphur dioxide gas: Cu + 2H2SO4 --> CuSO4 + SO2 + 2H2O You also get the following ionisation reaction with the water formed: H2O + H2SO4 --> H3O(+) + HSO4(-)

If you want to make hydrogen in a safe and inexpensive way, then I would go for dissolving aluminium foil in dilute hydrochloric acid (appr. 10% HCl works fine), to which a small amount of a copper salt is added. The following procedure makes hydrogen gas nicely and safely and requires no other chemicals than those available everywhere at low cost. 1) Take 100 ml of dilute hydrochloric acid (10% HCl), do not take the concentrated stuff. 2) Take two copper wires (electricity wire is OK) and perform electrolysis of the hydrochloric acid with these at a few volts. At the negative pole some hydrogen gas is formed, the positive pole dissolves. Don't bother getting the hydrogen gas of the negative pole, this is not why we want this electrolysis. We want this electrolysis because of the copper salt, which dissolves in the hydrochloric acid. Perform the electrolysis, until the liquid has a nice green color. 3) Add some aluminium foil to the green solution. You will have instant evolution of a lot of hydrogen gas and heating up of the liquid. If you have access to copper sulfate or copper chloride, then you can skip the electrolysis step. Then dissolve a few spatulas full of copper sulfate or copper chloride in the acid and then add the aluminium foil. ------------------------------------------------------------------------------------ I also know some alloys, which give instant hydrogen production, when added to water, but, although these are not expensive (speaking industrially), these by no means are OTC materials for individuals. These alloys are based on sodium metal and that is not a material you can easily find. The method, described above can be done by everyone.

Yes, you are right. You get hypochlorous acid. This is a very weak, but strongly oxidizing acid. This acid also is quite unstable. It slowly decomposes, giving chlorine, chloric acid and hydrochloric acid.

How do I know this? Because I read some books about this subject and I studied it . This is very basic chemistry. At the cathode you usually obtain hydrogen gas and at the anode the reaction products strongly depend on the electrode material and on the electrolyte (the dissolved salt). Metal anodes dissolve, giving the metal salt, except for the very noble metals and for niobium and tantalum. With graphite anodes (and also noble metal anodes like platinum, iridium, osmium) you obtain oxygen at the anode, or free halogen, when an halogenide is used as electrolyte.

Indeed, the Na(+) ions and Cl(-) ions are not oxidized or reduced. You can oxidize the Cl(-) ions when you use a graphite anode. The precipitate IS white, but sometimes it looks a little blue, just as cigarette smoke sometimes looks blue, while in reality is not blue. So, nothing changes with my story.

The gas you obtained definitely must be hydrogen. Probably you did something wrong with testing its flammability (letting the gas escape into the air, before you made it burn). The precipitate most likely is Al(OH)3. At the anode, aluminium dissolves, giving Al(3+) ions. With hydroxide, formed at the cathode, this forms Al(OH)3. The light blue stain you have also most likely is finely dispersed Al(OH)3, which is absorbed by the wood somewhat. The spilled liquid contained water, NaCl and Al(OH)3. It may have been absorbed somewhat by the wood and fine particles of Al(OH)3 may now be inside the wood. This is not a good thing. It will be hard to get it out. Using chemical means like dilute hydrochloric acid may help getting the insoluble Al(OH)3 out, but I'm not sure what it does with the wood and the laquer on the wood. To me, that does not sound as a good thing. Very finely dispersed white flocculent solid, such as the Al(OH)3 may have a somewhat bluish appearance. Think of cigarette smoke, which by no means is blue, but with certain light conditions, it looks blue. This effect also causes the very light blue color of the precipitate and the color of the stain on the table.