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I don't think so. H2 can perfectly exist in air, only at elevated temperatures it is ignited. The only thing, which could affect the reaction is the presence of dissolved oxygen in the acid. The oxygen may react with the Mg to form Mg(2+) ions and the acid reacts to water: 2Mg + O2 + 4H(+) ---> 2Mg(2+) + 2H2O But once hydrogen gas is formed, the reaction cannot be affected anymore.

I have done both reactions and none of these lead to an explosion. It may be true, when you heat several tonnes of the chemicals in a confined space, but it certainly does not apply to the test tube experiments you do. These MSDS's always give the most horrible scenario. You may remember the Toulouse disaster with NH4NO3. But that was with 20000 kilo of the chemical, in a metal container, in a violent fire.

Air does not affect the reaction. With, or without air, hydrogen gas is formed in both cases.

Usually there seems to be a change of color (e.g. if you put a copper coin in acid, it turns much more reddish), but that is due to dissolving of an oxide-layer, or more generally, a layer of impurities and tarnished metal. Most metals do tarnish in contact with air, especially if they are humid, or handled by humans and grease and dirt stick to the metal (e.g. coins, tools).

Sodium nitrite Copper (II) salts Chromium (III) salts Cobalt (II) salts Bromine Iodine Chlorine Red phosphorus Hydrogen peroxide Sulphuric acid With these chemicals you can do an amazing number of really interesting experiments. Transition metal chemistry is a very wide and interesting area of chemistry, also for the home chemist. Halogen chemistry also is a very interesting subject, with many interesting experiments, and many interesting compounds, which can be made.

This will be extremely difficult for the hobby chemist. Only the most brave and well equipped persons manage to make Na-metal at home. The safety issues are enormous. Using NaCl is more or less out of the question. Its melting point is way too high and it is near impossible to achieve that in a clean vessel for reasonable quantities for prolonged time, and perfectly isolated from air. Some people have made attempts to do this with NaOH instead of NaCl. The melting point of NaOH is much lower, but the use of this is certainly as dangerous as the use of molten NaCl. NaOH is corrosive and what do you think of NaOH at 350 C? Again, you need perfect isolation from air, and a corrosion resistant vessel (IIRC a nickel vessel could do the job). Some people over here at SFN suggested the use of NaNO3 as a low-melting salt, but that is insane. The nitrate ion is a good oxidizer and under such conditions you won't get Na-metal, but most likely nitrogen gas, maybe nitrite ion and maybe NO and N2O at the cathode, but certainly no Na-metal. Conclusion: Making Na at home is extremely difficult and requires extensive equipment. For me, it is beyond what I do, too dangerous.

I expect that H2O2 oxidizes any NO2 to HNO3 directly, so that would facilitate the reaction, but I'm not sure what happens at higher concentrations of HNO3. Sometimes, H2O2 also can work as reductor and then it could reduce HNO3 back to NO2, itself being oxidized to O2.

If you like the nordic regions of Europe, then do the following. I've done that and it is really amazing to see so many different places. I made a round trip as follows: The Netherlands Germany Denmark Sweden Finland (by boat to Turku) Norway (leaving Finland at Utsjoki) Finland (going back to the south) along the Russian border (Kuhmo, Joensuu, Lappeenranta) Russia (St. Petersburg) Estonia Latvia Lithuania Belarus Poland Germany The Netherlands This is a long trip (~10000 km) and it takes you several weeks, but it is great. I did the trip by car. You see many different things, nature, meet people of different cultures and see nice large cities (most natably Copenhagen, Stockholm, St. Petersburg, Tallin, Minsk, Warsaw, Berlin). Sleeping was done in tents, and that can be done for free in Scandinavia, Russia, and the Baltic states. In the other countries you best go to a camping place, or to a hostel. Belarus is somewhat spooky, you can skip that.

The halogens dissolve more easily in the alkane mix than in water. So, you extract them from water and they go into the alkane mix. The solutions only are dilute, so the density will hardly be affected by the presence of the halogen. The alkane mix has lower density than water and the layer on top is the alkane mix, with the halogen dissolved in it. Iodine also dissolves with a different color in the alkane mix, it is purple, while in aqueous solution it is brown. (Btw, iodine only dissolves VERY sparingly in water, when some iodide is added, it dissolves more easily.)

Actually, it is good that it is so smelly. Before the gas becomes really dangerous for you, the stench is unbearable, so you will try to get away, before it can do serious harm (or even kill you). Why does is smell like it does? I don't know. This question can be generalized. Many chemicals have smells, unique to those chemicals. The molecules of those chemicals trigger receptors and that causes the smell. The pungent part of a smell can be understood (stinging and burning sense). If a chemical is pungent, then it starts damaging the tissue. Some chemicals are odourless on their own, they only give a burning sensation. One of them is HCl. If you breathe this, then you smell a very nasty burning/stinging thing, but besides that it has no real "juice" or "taste". Ammonia, however, has a certain "juice" besides the stinging/pungent feeling. Chlorine gas also has a certain "juice" as part of its smell.

You need to add that after. If you add it before, the urea will destroy the NO2 already, before any acid is formed. That method is nice for removing small amounts of NO2 (somewhat yellow HNO3), but a dark brown/red or green/yellow heavily fuming liquid is not something you want to clean with carbamide (urea). The urea usually contains other impurities and with the urea you also add impurities. If only small amounts of NO2 are in the acid, then that is not that bad, but with large amounts of NO2 you need large amounts of urea.

Well, if you have a test tube, filled with pure chlorine, you can easily see the chlorine. A 1.5 cm (usual thickness of a test tube) layer of chlorine is not impressively green, but one should definitely see the color. Have a look at this, at the picture of the left test tube: http://woelen.scheikunde.net/science/chem/exps/clo2/index.html Not impressive, but visible. The gas ClO2 is much more impressive. However, electrolysis of NaCl (with graphite electrodes) does not give very pure chlorine at the anode. Especially if the applied voltage is too high, then a fairly strong side reaction occurs, where water is oxidized and oxygen + acid are produced at the anode, instead of chlorine. So, in practice you get Cl2, but also quite some O2. Things are even worse, if you collect the gas in an inverted test tube. Cl2 is soluble in water quite well (IIRC 3 liter of gas per liter of water), while O2 hardly is soluble (IIRC a few tens of ml per liter of water). So, a mix of Cl2 and O2 is even enriched more in oxygen, because the Cl2 dissolves in the water, and the O2 does not. So, your gas mix in the test tube may contain several tens of percent of oxygen. So, if you really want to make chlorine for an element collection, do as Jdurg did. Produce the gas by chemical means and clean the gas by bubbling it through water, followed by leading it through a drying agent. I would not prefer NaHCO3, because it will lead to contamination with CO2, I would use CaCl2 instead, which is very cheap and easy to obtain at hardware stores as drying agent for buildings. Also, I would not use Ca(ClO)2, but would use TCCA or Na-DCCA instead. Commercial Ca(ClO)2 contains up to several percents of CaCO3, and the chlorine gas may contain more than 20% of CO2. For most chemical experiments, this is not a real problem, but if you have a small sample for displaying, you want the gas as pure as possible, simply because its color already is quite faint, and diluting it with CO2 makes it even more faint.

Jdurg, I can only partly agree with you. Making HNO3 of not too high a concentration can be done with NO2. When NO2 dissolves in water, it disproportionates to nitrogen(III) and nitrogen(V). The nitrogen(III) in turn gives rise to formation of NO and NO2. Have you ever tried shaking NO2 with water? I did, in a sealed vessel. The deep brown gas quickly dissolves and what remains is a colorless liquid and a colorless gas. As soon as the gas is allowed to make contact with air, thick brown fumes are produced again. The reaction is: 2NO2 + 2H2O <---> HNO2 + HNO3 HNO3 <---> H(+) + NO3(-) (very strongly to the right, so HNO3 is taken away from the system) 2HNO2 <---> N2O3 + H2O N2O3 <---> NO + NO2 NO escapes as colorless gas, it hardly dissolves in water. So, the net reaction is: 3NO2 + H2O --> 2H(+) + 2NO3(-) + NO The NO can be reused, by allowing air to enter the mix. Then quick complete conversion to nitric acid occurs, because NO reacts easily with air to make NO2. -------------------------------- Now the part, why I still do partly agree with Jdurg. The story, given above, only holds for low concentrations of HNO3. As soon as the concentration of HNO3 rises, then the NO2 remains dissolved. The ionization of HNO3 is not complete anymore. NO2 remains in solution, N2O3 also remains in solution, the liquid becomes brown/yellow, or even green, due to the blue color of N2O3. So, when one continues bubbling NO2 through water for a long time, then one finally ends up with acid of a few tens of percent (30%, 40% I do not know exactly), with lots of dissolved NO2 and N2O3, which is heavily giving off red vapor on contact with air. So, this method of making nitric acid only is suitable for making dilute to moderately concentrated acid.

Para-dichlorobenzene was used for them where I live, but that practice is abandoned, because this chemical is believed to be a carcinogen. So, I definitely would NOT want to smell that too often. Almost all chlorinated organics are banned from the general public life, for a good reason, and p-dichlorobenzene is one of them. Naphtalene is the current alternative, but I also think it is not good to smell them. Btw, I hardly can imagine that you like that smell, it is terrible.P-chlorobenzene indeed has a peculiar, somewhat refreshing, non unpleasant smell, but the smell of naphtalene is plain crap. All chlorinates hydrocarbons I know (carbon tetrachloride, trichloro ethylene, chloroform, solventane, p-dichlorobenzene) have more or less pleasant smells, but sadly, all of them are quite toxic, and almost all of them are carcinogens and as such they are banned for general use.

Now the other part. Electrolysis of sulfate usually gives oxygen at the anode. But under certain conditions (relatively low pH, very high current density at anode), peroxosulfate can be formed. The mechanism is believed to be as follows: HOSO3(-) - e --> HOSO3• (• means radical) 2HOSO3• --> HOSO2-O-O-O2SOH The latter is peroxodisulphuric acid, H2S2O8, which is used to make persulfates, such as Na2S2O8 and K2S2O8 (these can be purchased as PCB-etchants, as alternative for the FeCl3-based etchants). The acid cannot be isolated, but salts of the acid, esp. the potassium-, sodium-, and ammonium salts are quite stable and are very strong, but somewhat sluggish, oxidizers.

NO NO!!! How many times this is written already???? Remember: SODIUM CANNOT BE MADE FROM AQUEOUS SOLUTION. Another time: SODIUM CANNOT BE MADE FROM AQUEOUS SOLUTION. Got it??? If you want to make sodium, then you have to work with molten salts, or maybe with mercury, making amalgams. But those are not the things to do at home.

I was the one who made that naughty post :D I still have the TACN in its vial, so it stores really well. It still can be made popping when put in fire, but real detonations I did not have. The CuO was just one way to het copper (II) nitrate in solution. So, I think it can also be done with copper, dissolved in nitric acid, but the complication is that NO2 is formed, and your solution will not only contain copper (II) nitrate, but NO2/N2O3 as well, which is not good for making TACN. So, if copper is used, instead of copper (II) oxide, then (1) Perform the reaction OUTSIDE, due to formation of copious amounts of NO and NO2. (2) After all copper is dissolved, boil the solution for a while, to drive off all remains of NO2/N2O3 from the solution. Use excess copper, such that not all of it dissolves, not even after heating. That gives best yields. Btw, I suddenly see that this is in a pyrotechnics thread. I would not really call this pyrotechnics, but that is another issue.

Measure the pH, e.g. by titrating with a solution of NaOH of known concentration. Especially if the concentration of acid is high, this gives a good and accurate result. Use an indicator, which changes color, when pH goes from 4 to 5 or somewhere around that. Using such an indicator, you take into account the somewhat acidic properties of Al(3+)-ions.

Whether or not a catalyst is used, the amount of energy by decomposing hydrogen peroxide to water and oxygen remains the same. Catalysts only affect the speed (the mechanism) at which reactions occur, but they don't add or remove energy. Hydrogen peroxide does not burn, it decomposes. The energy of decomposition of 2H2O2 --> 2H2O + O2 I don't know by head. Google definitely will give you the answer. It most likely will be specified as KJ/mol, but you can easily transform that to kJ/kg.

Sodium acetate and acetic acid are conjugate base and acid. The only thing they do is form an equilibrium. The mix of the two forms a buffering solution, with pH depending on concentration of the two chemicals. It makes no sense to titrate sodium acetate with acetic acid, nor the other way around.

It may explode, or it may work like a rocket for a very short time, but it definitely will not be a very good rocket. Normally, rocket fuel is solid or liquid. What you are talking about is a gas mixture. Density of gas mixtures is 100 ... 1000 times less than density of solid and liquids, so the amount of energy, available for driving the rocket upwards is very small. If it lifts, it will only do so for a fraction of a second.

This is not true. It is oxidized to Cl2. At the cathode, water is reduced to H2 and OH(-). Only by good mixing, the Cl2 and OH(-) react to form ClO(-). Here is problem #1. The mixing in practical setups can be a serious problem. Chlorine gas tends to bubble out of solution. Again this is not true. The hypochlorite disproportions to chlorate and chloride, but it only does at acceptable speed at elevated temperature. The hypochlorite does: 3ClO(-) ---> 2Cl(-) + ClO3(-) The Cl(-) then again can be oxidized to Cl2 at the anode. This reaction only happens with great difficulty. As long as there is some chloride left, the formation of chlorine is strongly favored, even when only a small amount of chloride is left. Besides that, formation of oxygen and acid at the anode becomes more and more a problem. Formation of perchlorate requires high current density at an inert anode. Even Pt-anodes are corroded severely under these conditions. A good anode is a RuO2-anode (usually Ti-metal covered by RuO2). The cathode is not that important. As long as a voltage is applied the cathode is protected, due to the strong negative potential at the cathode. Any cathode will do, but best results are obtained with cathodes, which also withstand the conditions of the cell for a some time, when no voltage is applied. A nickel cathode, or a titanium cathode are perfectly suitable, but if nothing else is available, even copper wire or stainless steel could do the trick, provided they are not allowed in contact with the liquid, when no voltage is applied. For the anode, titanium is totally useless. It erodes quickly and no chlorine is produced at all. You get a flocculent precipitate of hydrous TiO2 with that. For chlorate production, Pt is very nice, but graphite also could be used. Such small batteries are TOTALLY useless for chlorate production, let alone for perchlorate production. You need a sturdy power supply, capable of delivering plenty of current for prolonged times. Even at 7 A you only obtain at most 1 mol of KClO3 per 24 hours. Something which does work is a good lab power supply. You do not need such a fancy one with variable voltage and current regulation. A fixed voltage device can be used perfectly. If you don't want to spend much money, you could modify an old PC powersupply, often they can be obtained for free. http://woelen.scheikunde.net/science/chem/misc/psu.html If you don't want the hassle of modifying a PC powersupply, then something like this is suitable (you should be able to find something like that for around $50): http://www.rigpix.com/psu/propower_ps1310.htm It is not a state of the art power supply, but for the electrolysis it is perfectly suitable. With this, you also need a resistor network for current control, but that need not be a problem. For each mol of KClO3 you need at least 6 mols of electrons. For each mol of KClO4 you need at least 8 mols of electrons. In practice you need more, due to back-reduction at the cathode, where hypochlorite or chlorate is reduced back to chloride.

It is not simple at all. From a theoretical point of view, it is not just a matter of oxidizing the chloride anion to chlorate. Things are much more complicated through a multi-step reaction over chlorine, hydroxide, hypochlorite and finally chlorate. From a practical point of view, things also are not that easy. You need good electrode material, preferrably some way of current control (which can be as simple as a resistor network) and some means to prevent back-reduction at the cathode. All these subjects, I have covered in more detail on SFN, use the search engine and you'll find.

As with many things. Use common sense. There are risks, but they can be managed, but yes, you have to know the risks.

I agree with John, it is true for any manganese compound (remember, all manganese compounds easily are oxidized to either +2 or +4 oxidation state, depending on pH). On the other hand, if you have an occasional exposure, due to experimenting with MnO2, then in a practical sense I agree with YT. I myself sometimes experiment with manganese compounds and I take the usual precautions, but do not treat it in a special way as a highly toxic chemical. But if I had to work in a place with manganese around 5 days per week, then things would be another matter.