woelen

The acetic acid is strong enough to cause formation of nitrous acid. Nitrous acid (HNO2) is quite unstable. It has the following equilibrium reaction: 2HNO2 <---> N2O3 + H2O The compound N2O3 is blue. N2O3 decomposes very easily: N2O3 <---> NO + NO2 NO is bubbling out of solution. NO2 reacts with water, to form HNO3 and more NO (you are invited to derive the reaction equation yourself). If the temperature is too high, then the decomposition of N2O3 is going faster and NO is produced at higher speed, and the HNO2 hence disappears quickly.

Tartaglia, are you sure about the concentration of H(+)? A solution of a dichromate is somewhat acidic. Besides that, the reaction consumes quite a lot of H(+), so soon, at the very low concentrations of H(+) involved here, the concentration of H(+) rapidly changes. A cell, based on this quickly changes properties (deteriorates, because the reduction of chromium (VI) to chromium (III) becomes much harder at increasing pH).

Sodium- and potassium iodate are VERY similar with respect to energetic behavior. Just think of sodium chlorate and potassium chlorate, or sodium nitrate and potassium nitrate. Of course, the potassium salts are more suitable for pyro-work, but that is not because they are more energetic (in fact, they are even less, they deliver less KJ/mol of energy), but because they are not hygroscopic. And YT, I understand why you want to be very cautious. You have a lot to loose, just like I have. So, indeed, if you don't know the properties of a certain compound, then you'd better be very careful. But in the case of both KIO3 and NaIO3, trust me, they can be used for many experiments without the need to fear accidental explosions and so on.

YT, again, KIO3 is not shock sensitive in any practical sense. That MSDS is exaggerating strongly. I fully agree with encipher. I have done a lot of experiments with KIO3. I heat the stuff, mix it with all kinds of acids, and I molest the material in many other nasty ways , but it simply is not energetic enough to be of a real danger in those situations. If you dare to mix KClO3 with conc. HCl for etching PCB's, or if you dare to heat a solution of KClO3 in water, then you certainly can heat a solution of KIO3/NaIO3 in water. Iodate is to chlorate what a scootmobile is to a Ducati 999R.

Crystals of KNO3 are needle-like and crystals of KClO3 are like little scales. Yes, they are quite different.

YT, iodates are perfectly safe on storage and they are not shock sensitive. I have 250 grams of KIO3 lying around and this stuff is perfectly stable. It is an oxidizer, but its energetic properties are only a pale shadow of the energetic properties of KBrO3 and KClO3. Mixing KIO3 with reductors hardly gives an energetic mix. It just smoulders and smokes, but no violent reactions as with KBrO3 or KClO3. KIO3, however, can be used for very interesting experiments. It allows you to make KICl3, which is an energetic compound and shows lovely reactions, when mixed with finely powdered magnesium, such as giving purple clouds of iodine vapor on ignition (it ignites when mixed and a drop of water is added). http://woelen.scheikunde.net/science/chem/exps/KIO3+HCl/index.html Iodate probably does kill your plants. But, how much NaI you would like to add to your plants? You realize that iodine is a trace-element? It only is required in microgram quantities and having too much is quite harmful in the long run, both for the plants, but also for you, if you consume the plants.

YT, no need to worry about lots of lost chemicals. Add a VERY small amount of additional iodine, until your solution become light brown. Heat it a little to have your crystals redissolved again. Then proceed as described in the post above. @encipher: Making sodium iodate in the pure state is not that easy. It dissolves quite well in water and is hard to separate from the sodium iodide. You could, however, make potassium iodate fairly easily. Solubility of potassium iodate is much less. So, if you start from concentrated KOH and add iodine, you could precipitate KIO3 quite easily. This route to iodate from iodine, however, is not a very efficient route, because only 1/6-th part of the iodine goes into the iodate, the rest goes to iodide. Making iodates usually is done as follows: Take some iodine, mix this with solid KClO3. Add some concentrated nitric acid and start the reaction with careful slight heating. When this is done, then the reaction sets off and a lot of chlorine gas is released and a white solid is formed. This white solid is a mix of KIO3 and HIO3, forming an impure double salt KIO3.HIO3, with remains of nitric acid and chloride/chlorate trapped inside. This mix then in turn is separated from the liquid and dissolved in a small quantity of hot water and carefully KOH is added, until the liquid is slightly alkaline. Then the liquid is allowed to cool down and crystals of KIO3 settle at the bottom. These are fairly pure. Removal of the last traces of chlorate/chloride/nitrate can be done by redissolving this solid in an as small as possible amount of hot distilled water and letting crystallize again and slowly cooling down to almost 0 degrees C.

When you dissolve iodine in a solution of NaOH you do not only get sodium iodide! You also get sodium iodate: [ce]6OH^{-} + 3I2 -> IO3^{-} + 5I^{-} + 3H2O[/ce] Sodium iodate apparently is somewhat less soluble than sodium iodide (although it still is quite soluble). By dissolving iodine in a solution of an hydroxide salt it is hard to obtain pure iodide. The iodate, formed in the reaction, is a very disturbing impurity. ============================================================================== If you really want sodium iodide in a fairly pure state, then you could do as follows: Dissolve some iodine in a solution of sodium hydroxide. Use a slight excess of iodine, such that the solution is light brown. This solution contains iodide-ions, iodate-ions, sodium ions and a very small amount of free iodine (loosely bound to iodide-ions, but that is no problem). Next, evaporate the solution to dryness and then put the dry stuff in a crucible (porcelain, not metal, because that gives other impurities). Next, heat strongly. The tiny amount of iodine, remaining in the solid evaporates, the solid will become white. Also, the iodate decomposes, giving off oxygen, leaving iodide behind. Unfortunately, there most likely also will be a side reaction, where the sodium iodate (NaIO3) decomposes, not to give sodium iodide and oxygen, but iodine vapor, oxygen and Na2O. But I expect this impurity to be of very low concentration.

Plant ashes consist mainly of salts, which were in the plants. They are a rich source of potassium ions. A main constituent of plant ashes is potassium carbonate. Of course, many other salts will be in ashes, such as sulfates, chlorides, salts of sodium, calcium and magnesium. These salts make ashes appear light (grey, or even white). The grey/black color is caused by remains of soot, but also remains of soil can be present in the ashes. As you see, it is a fairly complex mixture. The salts can be easily extracted, by adding water to the ashes. Part of it dissolves, what remains solid is the carbon remains and remains of soil. When allowed to settle, a colorless liquid with the dissolved salts is formed above a grey/black precipitate. When this liquid is evaporated, then a nice white salt mix remains behind.

Dichromates are not so difficult to make in small quantities, and besides that, you do not need to isolate the dichromates. You have KNO3? And you have NaOH? If so, go to a ceramics supplier and buy some Cr2O3 (green chromium oxide). A few grams is sufficient. Melt a few spatulas of NaOH in a small inox spoon, add a spatula of KNO3, which also melts into the NaOH and then add some Cr2O3. The liquid will turn yellow. That is due to formation of chromate. Only use very small quantities, because working with the molten NaOH is quite dangerous. Let the yellow stuff solidify and add this to a small quantity of water. You'll get a yellow solution, probably with a lot of insoluble crap. Let settle and then decant the clear yellow liquid. This clear yellow liquid contains chromate (and still traces of nitrate). This yellow solution should be acidified with hydrochoric acid carefully, until the solution turns from yellow to orange yellow (not further, its pH should remain 7 or above). If the liquid becomes bright orange, then too much acid is added. Then ou need to add some NaOH again. This almost neutral liquid can be added to water, in which you dissolve your salt. One spatula full of Cr2O3 will give sufficient chromate for several 100's of grams of salt. Another way of making chromate is as follows: Get yourself some chrome alum (photo chemical suppliers, some drugstores also sell this, it is a totally non-suspect chemical, used in hardening fixers in photography), dissolve a spatula of this in water. Add a solution of NaOH, until the precipitate, which first forms, redissolves again and you get a clear green solution. Next, add some 3 to 5% H2O2 and carefully heat the liquid to 60 .. 70 C. The color will change from deep green to lemon-yellow (if too much chrome alum or H2O2 is used, the liquid may become somewhat brown/yellow, that does not really matter). Once the liquid is yellow, boil vigorously for some time to get rid of the last traces of H2O2 and then carefully add some hydrochloric acid to make it orange/yellow again (chromate is converted to dichromate). Then dilute with water and dissolve the salt into that liquid. Beware: chromate and dichromate contain hexavalent chromium and compunds, containing that probably are carcinogens, so be very careful not to be exposed to these compounds.

Conclusion: Chlorates are quite bad for pyro-applications anyway. Who can guarantee decent purity with chlorates, especially home-brew stuff? That is why [w00t] should be VERY careful with his home-brew chlorate anyways, regardless of what chemicals are used to make it. Chlorate is, however, a fun chemical to play with, in small demonstrations and redox experiments, for more serious work it is too dangerous. I myself hardly use any chlorate at all (although I have almost 1 pound of this stuff around). I do not do real pyro-stuff (not allowed in The Netherlands) and for aqueous redox chemistry chlorate is not the best oxidizer. It does not react cleanly, a lot of side reactions occur, when chlorate is used as oxidizer (formation of chlorine, chlorine oxide(s) besides the main reaction, where chloride is formed). This makes redox experiments, where one wants precise control over the reaction products almost impossible.

@YT: Why not mix chlorates and nitrates? But are oxidizers and as far as I know they do not react at all. to my opinion it is perfectly OK to add a concentrated hot solution of KNO3 to a concentrated hot solution with NaClO3. The KClO3 then will settle as nice crystals on cooling down. @[w00t]: Did you really buy 5 W rated power resistors? What size do your resistors have? Mine have a length of appr. 4 cm and a width of appr. 6 mm. They are the off-white ceramic things, rated for 5 W power consumption. However, if your resistors only become "hand-burning hot" then I would not worry too much. My resistors sometimes become VERY hot. But ceramic resistors can stand very high temperatures. Temps of 125 C are no problem at all and such a temp is not "hand-burning hot", but very hot, causing immediate blistering of your skin. Attaching a heat-sink to the resistors, however, is perfectly fine and if it makes you feel more comfortable, I would say, do so. If you have a few crystals of potassium dichromate at hand (for 100 grams of salt, 0.1 grams of K2Cr2O7 will do perfectly well), then you can increase the efficiency of your cell considerably. What happens is that chlorate ion is back-reduced at the cathode, when the concentration rises. What happens over there is the following: ClO3(-) + 6e + 6H2O --> Cl(-) + 6OH(-) besides the formation of hydroxgen. So, part of the chlorate present in solution is destroyed again at the cathode. This effect can almost be brought to an halt with just a tiny amount of dichromate present in the solution. Ask at your school if you may have 1 gram of K2Cr2O7. With that you can easily prepare 1 kilo of salt-solution. Beware, however, K2Cr2O7 is very poisonous and it is a probable carcinogen.

There are quite some chemicals, which instantly cause VERY severe burns on your skin. These chemicals, however, are not present in the average home, and not even in the average home chemistry hobby lab. Some of these chemicals are: - pure Br2 - Mn2O7 - HF - HNO3 (99%, not the standard 65%) - Cr2O2Cl2 Of course, there are even more extreme ones, such as SbF5, BrF3, MnF4, SO3. These are so corrosive, that they cannot be handled without special equipment. For this reason, such chemicals only are used in well-equipped labs.

The usual method is filing or drilling for medium fine powder. A chemical method, which produces ultrafine iron powder (pyrophoric!) is to make ferrous oxalate and heat this. The ferrous oxalate then decomposes, giving iron metal as a very fine powder, and carbon dioxide. Ferrous oxalate can be made by adding oxalic acid to a computed amount of Fe(OH)2, suspended in water. The solution, this obtained, is evaporated to dryness and what remains is ferrous oxalate. However, this procedure is not easy to perform at all outside a lab. You need perfectly oxygen-free operation. Fe(OH)2 is VERY prone to oxidation by air and during the heating of the ferrous oxalate, one also has to assure that absolutely no air (oxygen) is allowed to come in contact with the powder.

[woot], these however are quite dangerous. I'm afraid that there is no safe chemical, which ignites on contact with water, because this property on its own already is a very unsafe thing .

The computer does not crash, it goes on forever . A nice example is this loop for solving the equation x = cos(x) numerically: x = 0; do { x = cos(x); } while (true); It never stops. You need some stop criterion, e.g. x = 0; do { double c = cos(x); eps = fabs(c - x); // modulus of c - x x = c; } while (eps > 0.000001); This example is an iterative procedure. For summation of series, the stop criterion usually is that an error estimate does not exceed a given value. For alternating-sign convergent series, the error is less than the first term, skipped in the series. So, we can evaluate term by term, and when the term is smaller than the error threshold, then we stop. For same-sign convergent series, in general there is not a nice error-estimate. For each function, one has to to an analysis of the error as function of the index n, from where the series is truncated. Another REALLY important quirk when evaluating series for estimations of functions is that the smallest term needs to be summed first. Many small numbers may add up to a larger number. But, when the larger numbers are evaluated first, then the summing of the smaller numbers may be lost, due to finite precision of the computer. A somewhat exaggerated example, which clearly demonstrates what I mean: We have a computer, which works with four digits of precision, in floating point format. So we can represent numbers like 0.0000001, as 1.000*10^7. But because of limited precision, a sum like 0.1 + 0.0000001 equals 0.1 on this computer (remember, it only can store 4 digits of precision and an exponent). Now suppose we have a series with 1000 terms, each equal to 0.00001 There is one term, equal to 1. If we start with 1.000 in the total, and we add the smaller terms to the total, then the final outcome will be 1, because 1.000 + 0.00001 = 1.000 (remember, only 4 digits are maintained). Now, if we first add up the small terms, then we have 0.00001 + 0.00001 + .... (1000 times) = 0.01. Finally, we add 1.0 and the answer will be 1.010 You see? In computer science the following is not true in general: a+(b+c) = (a+b)+c. It only is approximately true. Because of this effect, programming the summing of series in computers is harder than one would expect at first glance. If programmed carelessly, one can easily loose a few digits of precision and that would be sad. A good software package determines all separate terms, stores them in an array of values, then sorts the values in increasing order of magnitude and then does the addition. Especially for large series this becomes very important!

There are many many good libraries, also as open source projects, which do the multi-precision math for you. Have a look at this: http://www.swox.com/gmp This is one of my favorite pieces of software and I have used it in quite some personal software projects, the most recent one being a generic polynomial equation solver for any type of roots and any required precision. In fact, high-precision maths works in quite the same way as plain good old schoolboys arithmetic, where the computer's natural number size serves as a single 'digit' and where addition, multiplication, subtraction and division are done in the way, children learn this at school. Only for VERY large numbers with thousands of digits, more advanced techniques are used for multiplication and division. One can resort to convolution or even FFT algorithms to perform multiplication in sub-quadratic time.

Some alcohols also can react with HCl or HBr, forming the chloro-substituted compound. E.g., refluxing isopropyl alcohol (rubbing alcohol) with conc. hydrobromic acid gives CH3CHBrCH3 and H2O.

Compilers and systems, which use fixed precision numerical data formats (such as IEEE 754) use stored constants. Also processors, such as the Pentium have hardwired bit-patterns for PI. There are, however, also software packages, which have user defined precision (e.g. computer algebra systems). These systems do not have stored values, but the values are computed as needed, but only once. The value is flagged as not available. When it is needed it is computed (which may take some time) and the flag is set. When the value is needed another time then its value is read again. As soon as the user changes the precision, then the flag of such constants is reset again, or the value is invalidated.

Measuring the chemical energy, contained in a combustible liquid or solid, in principle is easy. Burn a known ampount of the solid, and use it to heat some known amount of water. The water needs a specific amount of energy per unit of weight to be heated up (appr. 4.2 J/g/K). Using that, by measuring the increase of temperature, one can determine how many Joules of energy are contained in the material to be combusted. The practical setup, however, is not easy at all. A calorimeter should be used. It should be thermally isolated very well. One also has to take into account the heat capacity of the calorimeter (that device also will heat up). And even, if you could measure the amount of energy in the food, by means of combustion, that certainly is not the energy-content of that piece of food for the human body. The human body does not break down the food as far as during the combustion. Many cellulose-like compounds (fibers) are not digested and leave the body more or less unaltered, while at a combustion process these also are burnt. So, there are a lot of practical difficulties and uncertainties with this kind of measurements.

The main principle indeed is expressed by Yggdrasil. The combination oxygen/sugars has a higher energy content than water/carbon dioxide. This is how our body obtains the energy for its functioning. CO2 already is at a lower energy level and the body is not capable of obtaining energy from it. Plants can do things with CO2 (they use the C in it to build their own structures, mainly cellulose, but also sugars), but an external source of energy is required for them. That external source of energy is (sun)light.

Try this ! int a=10000,b,c=2800,d,e,f[2801],g;main(){for(;b-c;)f[b++]=a/5;for(;d=0,g=c*2;c-=14,printf("%.4d",e+d/a),e=d%a)for(b=c;d+=f[b]*a,f[b]=d%--g,d/=g--,--b;d*=b);}

At the temperatures involved, the CuO reacts with carbon to form Cu and CO. The CO2, which comes from the CuCO3 can also react with C to form 2CO. At such high temperatures, C/CO2 is less stable than 2CO. So, the net reaction can be CuCO3 + C --> Cu + CO + CO2, but it can also be CuCO3 + 2C --> Cu + 3CO The actual thuth most likely is somewhere in between and both reactions probably occur simultaneously.

Pure KClO3 burns really fast with charcoal, even if not mixed very well. If you light such a mix, then there is a 'whoosh'-like sound and it is gone. Btw, did you do the test with HCl?