Theophrastus

What is it that makes the reaction dangerous? The NO2 gas, I can account for with full glass apparatus, and teflon plumbers' tape if necessary, but what else? Is there perhaps a large amount of heat produced? Some other danger?

A little of topic, as this is meant to be a question of the rapidly evolving study of genetics, however, one of the pioneers of flight, a german inventor, Otto Lilienthal, designed ornithopters (flying machines that gain their thrust from flapping) which were actually relatively successful, despite the numerous dangers that the flaws of their design posed. I think that of all people on this earth, he probably enjoyed the closest thing to actually being a bird, as the ornithopters, were essentially giant systems of wings, attached to a person, however, these were incredibly dangers, that even their inventor could escape, as he was killed, after losing control, and going into a stall, falling back to the earth, lying eight hundred feet below. Shame...

Actually, I'm sad to say that the reason that people don't just boil down their peroxide to make rocket fuel quality concentrations, as that, the heat simply increases the peroxide's decomposition. It's likely that your resulting solution, based on how long you boiled it, remains at around 1-2%, due to the fact that just as the peroxide decreased, so did the amount of water. Sorry to spoil the fun...

Yeah, a further problem with astatine is that its half life, is just a little under 7 hours. Flourine is incredibly dangerous and reactive. In fact, its reactivity meant that it wasn't isolated until the late 19th century, and the chemist who did so, was awarded the nobel prize. If you look up incompatible compounds in the CRC handbook of chemistry and physics, you'll find the following to be true: Flourine incompatible with: everything ps: I agree, that you have to use all- glass apparatus when handling halogens. Certain types of rubber, based on the time of the experiment, and a low effusion rate, can also be used. Cheers

I was thinking that you could make the nitric acid, by decomposing a nitrate salt, by means of heat. This would form the metal oxide, as well as NO2. (Beforehand, I must stress that due to the dangers of NO2 gas, you must make sure that your apparatus is free of leaks. I generally use some kind of perfume, released in the main body of the apparatus, after which I wait several minutes, searching for an escaping scent, applying teflon plumber's tape if it is required (On occasion I've also made esters- purely for the fun of it of course: mmmm. pears.) Then NO2 is then bubbled through water, to form nitric acid. You can then store your newly made acid in a cool undisturbed place. As it is rather sensitive to light, use of tinted or opaque glass is recommended.

I meant it in a metallic sense; for example copper, is more noble than hydrogen, and as such, does not displace it. The ammonium ion, in a similar way, takes the same role as hydrogen, that of a cation, in binary compounds, and as such, would it not have its own place on the activity series? For example, were I to have on hand ammonium sulfate, adding magnesium, this displace the ammonium ion however, its precise place, I am unsure of. (NH4)2SO4 + Mg > 2NH3 + H2 + MgSO4

As many know from experience, single and double displacement reactions, are a common way of synthesising needed, and/ or interesting compounds. In this, a knowledge of the activity series is required, however, I have run into a problem: where on the activity series, would one place ammonia, or rather the ammonium ion? It was is one of few nonmetal cations, and as such, would belong on the metal activity series, but where? Any thoughts?

Based upon the size of your sample, I know that the corrosion process speeds up greatly in a pure oxygen atmosphere, so if you place your metal in a large flask of oxygen, you should quickly attain the necessary corrosion. I agree with using peroxide, or some other oxidiser, as in theory, it ought to work quite well, though I, myself have never tried it. Cheers!

I agree with hermanntrude, obtaining it from magnesium pencil sharpeners would probably be best. You can tell easily based upon the weight. Obtaining it through chemical means however, is rather complex, so the simplicity of this idea in contrast, makes it far better. If you require a strip, to use as a fuse, the metal is actually rather soft, and malleable. Certainly not as easy to cut as sodium, but certainly not such great an effort as steel. Best of luck, ,Theo

Really? Cool! I've never really done anything with it, and know it far worse than most other common acids. Interesting...

I know: safety first. Yes I did wear gloves, lovely chemical resistant ones I recently purchased from an online supplier, despite this, I still thought I felt something. Strange, though as I stated in my first post- irrelevant to the matter at hand. I think that an occurance such as this would better fit into the psychology forum page. Regardless... Cheers!!!

Thanks for the muriatic acid tip. I recently consulted, a friend, only to discover large 34% buckets, being sold at home hardware. Absolutely brilliant stuff. An interesting but unrelated thought: on first handling the mixture, I felt a slight stinging sensation, however, I rinsed my hands thoroughly to be sure, and in retrospect, it was probably only a figment of my mind's early anxiety, just as one feels as if their head is itching, at a checkup for lice, despite the pests' inexistence.

First: Were you running a current through the wires? (sorry for asking a crucial but stupid question) However I tried this same experiment quite recently in water, and obtained the same product. Second:What was the concentration of the peroxide solution, and why bother to do an electrolysis (essentially) When the oxygen formed by the peroxide will quickly be absorbed by hydrogen ions. Third: What was the concentration of the peroxide? General Thoughts?

While it may seem hilarious to the author in retrospect, I do not think that this is an experiment I desire to replicate. Personally, I recently read a book of a similar nature, absolutely brilliant, Oliver Sacks', "Uncle Tungsten: memories of a chemistry boyhood" It really quite a wonderful book. Interesting quote, however, is there any underlying truth to this, as it is true as the author states, that theoretically nitric acid does not react with pure copper. Perhaps it was some manner of oxidised layer of oxide, carbonate, hydroxide etc, and I do know that many coins, currently, only have a thin electroplated layer of copper on top, under which rests a cheaper metal, or metal alloy, generally of iron, nickel and zinc, in various proportions. Any thoughts?

A rather simple way to make copper chloride, is to start with copper oxide, which can easily be obtained by heating of pure copper in air (I recommend use of the powdered metal, if you have some, for this). While copper does not generally react with most weaker strong acids (I know, an oxymoron), copper oxide, you can easily react with hydrochloric acid, to produce Copper chloride, as shown below. CuO + 2HCl > CuCl2 + H2O In regards to obtaining copper oxide, you can also heat copper hydroxide, and this will decompose to form copper(II)oxide and water at a relatively low temperature, or if you have a stronger burner, perhaps even a bunsen burner, you can use copper carbonate as well, decomposing to form copper(II)oxide and CO2. Hope this helps!

In regards to wondrous flaming experiments, I think I have your fix. The simplest I know is the fire worm, where you mix baking soda, and sugar, then lightly sprinkle the mixture with some ethanol, leaving it to dry. You can then contain the reaction using a bucket of sand, poking small holes where you insert your solution, or even doe it right on a petri dish. You then ignite the the mixture, due to the heat, the baking soda (NaHCO3) decomposes into sodium carbonate, carbon dioxide and water. 2NaHCO3 > Na2CO3 + CO2 +H2O The heat will and pressure of the sugar will mold it, and your mixture will curve upwards, and as the reaction stops, a dark worm remains (rather small though) You can always blow up, a couple of hydrogen balloons, by using a match on a (preferably long) stick. You can quickly produce a reasonable quantity of hydrogen gas, by reacting aluminum foil with HCl or NaOH. You can also heat magnesium (be careful, as this will produce some pretty white sparks), then drop it into water. The equations for both reactions, are found below: 2Al + 6HCl > 2AlCl3 + 3H2 Mg + 2H2O > 2MgO + 2H2 You can also use various chemicals to colour flame; copper compounds such as copper sulfate and chloride, generally burn with a bright greenish blue, or blue flame, strontium and calcium salts are red, etc. There's plenty of choice! I personally don't have so great a fondness for fires and bangs, but I do enjoy complexes, what with their brilliant colours. These can be synthesised by mixing ammonia (a Lewis base) with copper compounds like copper hydroxide and sulfate. Copper hydroxide can be easily attained by electrolysis. This can then be heated to achieve copper oxide which you can either react with hydrochloric or sulfuric acid, to produce copper chlorides and sulfates. The same can be done with most hydroxides. CuO + 2HCl > CuCl2 + H2O CuO + H2SO4 > H2O + CuSO4 Have a blast! ,Theophrastus

Ok, to all of those out there who claim that it is easier to simply make things from scratch, rather than recycle them, I must say that this varies from product to product. Recycling metals, for example, requires far less energy, than is required to refine them from their respective ores, in the first place. Also, wide- scale recycling of a particular material is simple proof that it is simpler, as large- scale production on various fronts is a sign that a process is economic. For example, the reason that some alternative "green" products cost more than their wasteful counterparts, is due to the fact that as the energy required to produce them is greater than that of the common product, on the market (ie denim insulation vs. ordinary insulation) the price has to be increased, to achieve an adequate financial profit. Recycling works in many cases, and seriously speaking: Yes, everything you put in your blue and compost bin, is in fact recycled, as is the nature of recycling. Alas...human ignorance. Though I do agree, the original post in truth, is not a question, or even relevant as a discussion topic. If you desired a view on recycling, please, do post a note somewhere to display it to be so. Go green.

For your first question, compact flourescent bulbs are in fact more energy efficient (theoretically speaking; different bulbs, different brand names, etc.), generally speaking, as unlike incandescent bulbs, much less heat is lost, as the reason that incandescent bulbs light up, is that their tungsten filament provides a great enough amount of resistance to induce a glowing in the filament, however this resistance also results in electrical energy, being released as excess heat, whereas fluorescent bulbs use noble gases (primarily neon) which when electricity excites these atoms, spurns them to emit a photon. (light particle) Based upon the gas used, the spectra and wavelength of the light will differ, but that is the basis. While fluorescent bulbs are more energy efficient, and have a longer life expectancy (as the excess heat of an incandescent bulb results eventually in the filament "burning out"), they have been recieving a rather large amount of criticism due to the fact that they contain harmful mercury vapour which is very difficult to safely dispose of afterwards. L.E. D.'s however I'm not so familiar with, and that, I suppose is a question, someone more experienced can answer. Hope this helps!

seriously speaking, I doubt it would really matter, th reactivities of the two, as regardless, the ammonia, along with the sulfate, (The ammonium sulfate) will act as a ligand, forming a coordination complex with the copper.

Cool, though with ammonium persulfate, you'ld have to be careful that it doesn't form an ammonia- copper sulfate complex, following the ammonia's displacement.

Finding the lovely, yet rather old thread, you guys have on making copper sulfate, (which actually yields copper hydroxide), I was curious as copper hydroxide is a useful chemical to have around, and as such performed the experiment as prescribed, using copper electrodes, and magnesium sulfate (epsom salt) as an electrolyte. I used about 18 volts initially, however, as batteries burned out, some were replaced, some weren't, I was left with about 12 volts, by the experiment's end, however the results, were not at all as I expected, as the solution turned a soft , pale shade of yellow, while the precipitate was orange. I thought this to be due to an excess of epsom salt, so I tested its solubility, and appropriately decreased its concentration upon my second attempt, however the yield was the same. Since there was naught but epsom salt, I think I can make the assumption, that the orange precipitate is copper. (this I shall test by heating it, when I get some more denatured alcohol, for my alcohol burner next week) But what happened. Many people seem to have used this method, with excellent results, however, why didn't my method work? For those unfamiliar with the forum in question, the theory is that the copper "excited" by the electricity will displace the magnesium momentarilly. In an aqueous solution, the displaced magnesium would then react with the water to produce magnesium hydroxide. The magnesium hydroxide would then react with the copper sulfate, the magnesium displacing the copper to once again attain magnesium sulfate and copper (II) hydroxide: Cu* + MgSO4 > CuSO4 + Mg Mg + 2H2O> Mg(OH)2 + H2 Mg(OH)2 + CuSO4 > MgSO4 + Cu(OH)2 Thoughts?

Theoretically speaking, reacting nitric acid, with glycerine (a common by product of the soap- making industry (thank you fight club!)) will in fact yield nitroglycerine, however, I hope you only ask this in a conversational manner, because, I don't feel I even need to mention the volatile nature of this compound. Despite this, following its synthesis, it was widely used in various industry, but at the toll of hundreds of thousands of lives. Its very creator, ascanio sobrero, warned of this. It was only years later that alfred nobel, found a way to stabilize it, yielding the more stable, however equally deadly, dynamite. "(Which is a store for all you non- Americans)"??? Don't worry, I'm sad to say that Wal- mart is a global, not a national evil, that has yet managed to permeate, numerous nations. Alas...

Reading over an old thread, I found a fact that I haven't really thought of juxtaposing in to my vast array of ideas for synthesis experiments. Knowing that hydrogen is more active a "metal" than copper, could one not perform a displacement reaction forming sulphuric acid, by bubbling hydrogen gas, in a solution of copper sulfate, the result would then be sulfuric acid (as I before mentioned) and a copper precipitate. Would such a process work? In theory yes, but I have several doubts due to practical issues such as hydrogens lightness. Thoughts?

Ah, alas, a method gone wrong. I suppose upon buring, it would oxidize to magnesium oxide, reversing the reaction that originally took place. Shame... (Though thanks for the clarification)

Ah, perfect! Thanks for the input. Thus, it can be safe to conclude, that it was mostly magnesium oxide, caused by the reaction between Mg and NaOH, as so, I believe: 2Mg + 2NaOH > 2Na + 2MgO + H2 That explains the violent bubbling! as such, the vinegar would then react with it, to form magnesium acetate, releasing water, in the process. That, said, this implies that both precipitates would have to come off of the cathode, not the anode. The hydroxide, used up, there is only Copper, Oxygen and Chlorine, to speak of. As such I believe the lower precipitate is copper oxide, while the upper one... well that I'm unsure of. The only way it could be copper chloride, is if the solution is already saturated with CuCl2, however, as the solution remains clear, this is not the case. On an alternative note, can you think of any way to separate the two precipitates, from one another; I'm clueless. If I were to boil it off, would they not crystalize upon the beaker, and then, with like, clinging to like, could I not extract the crystals separately. Cleaning off any excess, stuck to the walls of the beaker, with acetic acid solution, or perhaps, a heavy- duty drain cleaner.