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Is your dad concerned about the creditcard (fraud and so on) or doesn't he want you to play around with (potentially dangerous) chems? If the first is the case, then there are other ways to get ahold of chems, with the help of your dad (eBay indeed is a good source). If the second is the case, then you have bad luck and you'd better wait for a few years or try to explain what you really want with the chems.

The fact that there is 'NO2' in it does not mean that it is brown. This way of reasoning does not hold. For example, NaNO2 contains 'NO2', but not as molecule, but as ion. The compound is not brown, but white/colorless or very pale yellow if finely divided. Another example is nitromethane, which contains 'NO2' as nitrogroup: H3C-NO2. This is a colorless liquid, like water. There also is an isomer H3C-O-NO, it also contains 'NO2', but now one of the O-atoms is between the C and the N. This is the colorless gas methylnitrite, which I have prepared some time ago from acidified methanol and sodium nitrite. You see, the fact that a certain group or set of atoms is in a molecule does not say anything about the color. If nitronate ion is brown, then it is a pure coincidence that this seems to match with the color of NO2. So, anyone out there, who knows the real answer?

I'm playing around with salts of the nitronate ion. All experiments I do point in the direction of a dark brown/red ion, but from a chemistry expert, I understood that nitronate ion is colorless. The expert, however, does not have personal experience with nitronate salts. I prepared the nitronate salt by adding a solution of KOH to a solution of nitromethane. Nitromethane, when dissolved in water, is converted to nitronic acid (the so-called aci-form of nitromethane): H3C-NO2 <----> H2C=N(O)OH The H-atom at the right is the acidic one. When hydroxide is present in solution, then the liquid turns red/brown. According to literature the following reaction occurs: H2C=N(O)OH + OH(-) ----> H2C=NO2(-) + H2O I originally thought that nitronate ion is colorless, but my experiments show a deep red/brown color. The solid can be isolated as a dark brown water-soluble powder. The reagents I use are very pure (99.9+% nitromethane and reagent grade KOH in distilled water). Still I get the dark brown stuff. So my question is: is nitronate brown? or is there a side reaction, which produces a brown compound?

Copper hyroxide in fact is quite interesting stuff. If you have tartaric acid (you can buy that at every drugstore), then you can make Fehlings reagent, which is very nice for making red copper (I) oxide or even plating thin solid layers of copper (I) oxide. You can also make the tetrammine complex with copper (II) hydroxide. Add an excess amount of ammonia to a suspension of Cu(OH)2. You'll see that it dissolves and that you get a really beautiful deep blue color. If you also have sulphuric acid (dilute is OK), then you can make the solid salt Cu(NH3)4SO4.H2O. If you have dilute sulphuric acid, then you can make copper sulfate crystals. These are very neat, bright blue. Add dilute sulphuric acid to copper hydroxide, until all of it has just dissolved and then boil away most of the water. Do not try this with nitric acid though. With nitric acid you get Cu(NH3)4(NO3)2, which is very explosive and having a gram or so of this may be enough to blow off your fingers!

You do not need to be afraid of an explosion. If you want to dissolve gold, then go for NaNO3 or KNO3 in HCl. The ammonium ions spoil the solution quite a lot, because a lot of active ingredient is used up for oxidizing the ammonium ion to N2 and N2O. If you don't have KNO3 or NaNO3, then you could try with ammonium nitrate, but it does not work as good as the other two. The compound ON-Cl is called nitrosyl chloride.

You definitely should keep some of the tin chloride for other experiments. Tin (II) forms many interesting complexes and precipitates. The most interesing one I have made so far is the deep orange/red [Pt(SnCl3)5](3-). It is so remarkable because of the SnCl3(-) ligands and besides that, it also is penta-coordinated, which is really special. But also if you do not have platinum-salts, there are many interesting experiments with it. E.g. add a very concentrated drop of solution of KI to a very concentrated drop of solution of SnCl2 and see what happens. You get a yellow precipitate, which after some time turns bright orange/red and also becomes dry and crystalline! Tin (II) in general is an interesting fairly strong reductor. BTW: Many concentrated solutions may indeed have a look as if some detergent is in them. HCl also is one of them. My 30% HCl keeps bubbles on the surface of the liquid for a long time, when it is shaken.

In the cold, none of these will lead to a reaction. With some heating you'll end up with an incredibly complex mix (with both of them). With heating you get the following reaction products: NO (colorless gas, brown in contact with air) N2O (colorless gas, this is due to the presence of ammonium ions) NO2 (brown gas) ClNO (orange gas) N2 (colorless gas, this is due to presence of ammonium ions) Cl2 (green gas, but only when no ammonium ion is present) The liquid turns yellow, due to ClNO, dissolved in the aqueous liquid. If you do the same experiment by dissolving KNO3 in hot concentrated HCl, then you get the same stuff, but with the N2O and N2 replaced by Cl2. Just try it at a test tube scale. Interesting reaction. Be careful though, the liquid is VERY corrosive and the fumes are a poisonous coctail! The liquid easily dissolves gold, especially when warm.

The blue precipitate at the cathode can be explained easily. At the cathode, the following reaction occurs: 2H2O + 2e ---> 2OH(-) + H2 The 2e is coming from the cathode, delivered by the power supply. The hydroxide ions form a precipitate Cu(OH)2 with copper ions, already present in solution. You use a very high voltage. If you lower the voltage, then I certainly expect less Cu(OH)2 to be formed and more Cu-metal. Another thing is that it indeed is not very good to use CuCl2. It is much better to use CuSO4. CuCl2 forms many complex ions in solution, such as CuCl4(2-) and other less-chlorinated complex species (these complex ions make a solution of CuCl2 appear green instead of blue). These complex ions also makes things more difficult.

I'm surprised to read that HCl is not that common in the UK. Over here you can buy it almost everywhere, in 10% and in 30% concentrations. @YT: Detergents? Did you means that is in some of your chems, or is it just looking as if there are some detergents in your chems. If you have 20+ % HCl, then you have decent stuff. It is of moderate concentration, but enough for many experiments. Is that liquid colorless or somewhat green? I have a bottle of 38% HCl. That is a pain to store. The bottle is somewhat pressurized, like a bottle of carbonated soft drink. If I open the screw cap, then I hear a hissing noise like you hear on opening a fresh bottle of coca cola. The pressurized HCl then comes out of the bottle giving a thick white choking fume. For my normal experiments I use 30% HCl though. That also gives fumes, but at least it is not under pressure. It is sold in 5 liter jerrycans. We also have 10% HCl over here in hardware stores, that stuff is of high quality but relatively expensive. No colored impurities at all and on evaporation hardly any residue is left.

Chlorate can be converted to perchlorate by means of electrolysis, but this must be done with the sodium chlorate solution. Potassium perchlorate simply is not sufficiently soluble and it will be really hard to get perchlorate, free of chlorate if you use potassium chlorate as starting point. I think you need an anode, made of sintered PbO2. For more info see the site of this Dutch guy: http://www.wfvisser.dds.nl/EN/kclox_EN.html It is not really easy, but with some effort and patience, it should be possible.

If you electrolyse KNO3, then do not expect to get any K-metal. At the anode you will get NO2 and O2, but the formation of the NO2 is just the smallest of your problems. At the cathode, if any K-metal is formed at all, you'll get an immediate redox-reaction between the nitrate ions and the K-metal. Remember, nitrate is a good oxidizer, used in pyrotechnics and it easily oxidizes weaker reductors, like sulphur, diverse metal powders, etc. So, imagine what happens if molten K is mixed with molten KNO3. You'll also get some KNO2, due to decomposition of KNO3 in the molten state, but I expect that effect will not be too severe, just above the melting point of KNO3.

Indeed, freezing out some ice and keeping the liquid is a fairly good way of concentrating. If you start with 100 ml and you allow 60 to 70 ml of to freeze, then the remaining liquid has somewhere between 6 and 8% of H2O2. I did this once and it works fairly well. It is said you can go up to 20 or 30% with this method, when repeated a few times, but I have no experience with this. However, as others stated, going beyond approximately 30% is not a wise thing to do. It is incredibly dangerous.

Indeed, I agree with you, situations just below the 'horizon' are not extreme at all. I understand your word 'tidal' , I think it also is the English word, in Dutch it is called 'getijde'.

Obtaining SO3 even is more difficult. Where making H2SO4 is quite hard for the home chemist' date=' making SO3 is next to impossible for the home chemist. On sciencemadness there are some people working on that, but it is incredible difficult, the stuff is incredibly corrosive and concentrated H2OS4 is just a children's toy, compared to pure SO3. See this link. If you add SO3 to water, then it detonates! In industry this is not the way of making H2SO4! They add the SO3 to H2SO4, such that it is diluted and then they add water to the SO3/H2SO4-mix (called oleum). As you can see, making H2SO4 and SO3 is not easy at all. Fortunately, H2SO4 is fairly easy to get your hand on. Getting SO3 is not easy at all, and I must say, that is for good reasons! Having lots of SO3/oleum around would be a MAJOR risk for the average home chemist.

Of course that does not tell you very much. A car also is not hygroscopic, but on a car, the next day you can see a lot of droplets of water on a cold and bright morning. So, even on a non-hygroscopic compound like KClO3 you can have some 'dew-like' settling of water. The difference with the car is that KClO3 is somewhat soluble, even in the cold, and that makes the material stickly and clumpy. If you, however, dry the material in a warm, dry room, then you'll end up with a perfectly dry powder or crystals. Another reason may be that the KClO3 contains hygroscopic impurities.

NaOCl(s)?????? Are you sure? I hardly believe this, because this stuff is unstable in the solid state, as far as I know. Isn't the stuff Ca(ClO)2 or Ca(ClO)Cl? These are stable solids. Hypochlorite can be used as a strong oxidizer in alkaline environments, it can also be used for making chlorine gas by adding some hydrochloric acid, but that should be done with the utmost care and certainly not inside. The chlorine gas may kill you, if you are not doing this in a careful way. For the careful user, however, chlorine gas is a very nice chemical to experiment with (e.g. try burning a stick of wood, immersed in liquid candle-wax or candle-stearin in pure chlorine, this is really neat to observe). There are many more interesting experiments with chlorine gas, just google around and you'll find quite a lot of them.

What do you mean with 'density' for a black hole. To me, this concept hardly makes any sense, because near the Schwarzschild radius, you just have empty space and near the singularity you have very weird conditions. Indeed, if you were to fall inside a large black hole, then at first you would not notice, but the closer you fall towards the singularity, the more you're torn apart (the difference in gravitational pull between your feet and your head will become larger and larger, because your feet are somewhat closer to the singularity, assuming you are falling inwards with your feet 'below'). The concept of density, as you mention, isn't that 'average density' over the entire volume of the black hole?

Dissolve as much as possible of the impure KClO3 in very hot water (almost boiling). Keep the liquid very hot, but not boiling and let insoluble matter settle at the bottom. Next, decant the hot and clear liquid from any insoluble impurities, which settled at the bottom and let the liquid cool down to room temperature. Crystals of quite pure KClO3 settle at the bottom. Other impurities remain in solution. If the liquid is at room temperature, then put it in a fridge for separating more of the crystals. Decant the liquid and keep the crystals of KClO3 and rinse them briefly with ice cold distilled water and then let them dry at air (may take many hours, but finally they will perfectly dry, because KClO3 is not hygroscopic at all).

Have a look at this. Lots of information on H2SO4. It definitely is not that easy to prepare in your kitchen. Be careful! http://www.sciencemadness.org/talk/viewthread.php?tid=510

Copper powder can well be stored without oxidation, but you have to be sure that the copper is not stored in an acidic liquid. So, rinse with some HCl and then rinse with a lot of water. Finally, rinse with a dilute solution of NaHCO3 (this removes any traces of acid, without itself being corrosive) and finally rinse with distilled water and let dry. In acid, you have the following reaction, due to oxidation by oxygen from the air: 2Cu + O2 + 4H(+) ---> 2Cu(2+) + 2H2O

As someone else already stated, plasma contains electrons. The electrons are not bound to a specific atom (just as part of the electrons are free to move around through a metal). Still, the Pauli principle holds and the nuclei and electrons cannot occupy the same space. So, you can have a conductive plasma with a density of severals tens of grams per cubic centimeter, but at a certain gravitational pressure the Pauli contra-force collapses and the matter inside the star degenerates to a srt of neutrons.

The Pauli exclusion principle assures that we have elements, as we know them. This is the reason, why solids and liquids for instance are not compressed easily. The electrons and nuclei cannot be at the same location at the same time (very roughly speaking ), orbitals only can accomodate two electrons with opposite spin. If the pressure on a solid is increased further and further, then at a certain point, the force, excerted on the electrons and nuclei becomes so high that it would break the Pauli exclusion principle. However, as that is not possible, another thing happens. The electrons and nuclei 'melt' together, forming neutrons. These neutrons occupy much less space and the matter can collapse incredibly (the density can increase by a factor of 10^10 or something like that). In fact, no new elements are formed, but the (neutral) atoms simply collapse, as soon as the back-force from the Pauli principle breaks. Of course, it might be that there are fractions of elements with partially collapsed nuclei and electron shells for a small moment of time, which may melt together to something which may be element-like, but this is just theoretical. Such entities quickly will 'melt' down with the remaining electrons. Now, in the neutron star the same principle holds. Neutrons also satisfy the Pauli exclusion principle and it is this again, which assures that the neutron star does not collapse further. However, when the mass of the neutron star also exceeds a certain value, then the pressure becomes so high, that the Pauli exclusion principle cannot hold anymore. This causes the neutrons to 'melt' together and at this point there is no (known) particle anymore, which satisfies the Pauli principle. So, the collapse continues forever and the density grows to infinity.

Both the dark stuff and the white stuff appeared at the SAME electrode' date=' being the anode. The white stuff is due to formation of SnCl2, which immediately is precipitated by the presence of OH(-) ions in the solution. The black stuff most likely is the lead, contained in the solder wire. The tin dissolves and the lead remains as a porous solid, which falls from the anode. The cathode hardly is attacked, it looses some of its shiny appearance. Indeed, no chlorine gas. The anode material is more easily oxidized than the chloride ions in solution. No, not use less salt, but indeed lower your voltage. The tin ions are reduced more easily than water molecules. If the voltage is too high, then hydrogen gas is formed, together with tin metal, but this makes it hard to get a nice cohesive piece of metal, due to the disrupting bubbling. With a lower voltage, only the metal is plated out. This, however, only works nicely, if the liquid is clear and slightly acidic, otherwise you don't get a nice clean metal, but some impure black crap. The HCl is needed to dissolve the Sn(OH)2 and making the liquid nice and clear again. Without the HCl, you may be able to plate out the metal, but you'll also include lots of other insoluble things. This is the reason that you don't get a nice shiny metal at the cathode, but some black gunk instead. .