woelen

With such a current it must take a LOT of time before you have an acceptable amount of barium chlorate. At 0.33 amps you have 0.33 Coulomb of charge per second. One mole of electrons is approximately 96500 Coulombs. So, for one mole of electrons, your cell needs to run approximately 3*96500 seconds. For making one mole of Ba(ClO3)2, you need 12 moles of electrons (each chloride converted to chlorate takes 6 electrons), so your total time needed for making one mole of Ba(ClO3)2 is 12*3*96500 seconds. This is approximately 40 days. The molecular weight of Ba(ClO3)2 is 304.23. This means that with your cell you'll get approximately 0.75 grams per day (24 h). Isn't that very slow? Probably we must use a whole array of anodes in parallel, with current evenly divided over each anode. Another possibility is to place cells in series. EDIT: The computation is wrong, 0.75 must be 7.5 With that result the conclusion is not valid anymore, see response further below the thread .

Take two little plastic caps. One cap should be filled with 1 ml of conc. hydrochloric acid and the other cap with ammonia solution. Put them next to each other or even better, put the hydrochloric acid cap just behind the ammonia cap, such that the ammonia vapor blows over the hydrochloric acid. This gives a LOT of white smoke, which is totally safe and you probably will have some excess ammonia vapor (this is released somewhat quicker than HCl vapor). Ammonium chloride is non-toxic (in fact, it is used in certain sweets, called salmiac) and ammonia also is not toxic in these amounts. You have to be careful though that the caps cannot tumble and that their contents is spilled on the device to be tested in the wind tunnel. Conc. hydrochloric acid is very corrosive, especially towards metals.

That is good to read. This perfectly is in line with my observations. As soon as I get oxygen production in my bromate or iodate cell, the anode starts to pulverize. Before that, all is working fine without the pulverization.

Good to hear that a current source is better, so my idea was not that bad . A current source is no problem for me to make. I'll see what scrap components I have around. I can make a current source with any opamp. Btw, you do not have to worry and need to add a 1 Ohm resistor in the output line of a current source. That is the nice thing of current sources. When the output is shortcircuited, it simply gives off the current for which it is designed, so having a shortcircuit does no harm when a current source is used. Having an open circuit can do more harm with a badly designed current source, because that will cause the output voltage to be driven high and the components do not operate in their normal mode of operation.

I have done quite some electrolysis experiments of halogenide solutions (NaCl, NaBr, NaI) with graphite anodes and i have observed the following things: 1) At low voltage, the halogen is mainly formed at the anode. With bromide and iodide, the electrolysis can be totally free of bubbling oxygen from the anode. 2) At higher voltage, a side reaction occurs, in which oxygen is formed. When this occurs, there also is considerable pulverization of the anode. 3) After a long time of electrolysis and mixing, also more oxygen is formed at the anode, also with bromide and iodide. This formation of oxygen is associated with pulverization of the anode. 4) The current depends on the voltage in a highly non-linear way. Above a certain threshold voltage, the current drawn by the circuit increases sharply with increasing voltage. This threshold voltage is highest for chloride and lowest for iodide. I have an idea to use a current source instead of a voltage source. It is easy to make a current source of 1 A with a simple OPAMP and a moderate power transistor. Using standard graphite rods from batteries with a diameter of 7 or 8 mm and a current of 1 A one should not exceed a current density of 0.1 A/cm2, when the electrode is dipped into the liquid for at least 5 cm. This should prevent strong pulverization, according to Raivo. Using a current source has the advantage, that it automatically adjusts at the halogen to be electrolysed. For an iodide solution a lower voltage is 'chosen' by the current source than for a chloride solution. This eliminates the problem of using too high a voltage, resulting in unwanted formation of oxygen and associated pulverization of the anode. When the electrolysis comes to an end and the amount of halogenide is getting low (i.e. lots of halogenate) in solution, then the current source automatically increases voltage in order to keep the current at 1 A. If there is someone with comments on this, I would appreciate that very much. I'm looking for ways to do electrolysis in a more controlled way for many different electrolytes and I have the idea that a current source is better than a voltage source.

Ah, that is interesting info... So, current density should not be too high. Your response gives me some interesting idea's. I'll start a new thread on electrolysis, here we go too far off-topic.... Any responses on electrolysis, please go to the new thread, which will appear within a few minutes. Here please continue on the KBrO3 and KIO3 and related things .

Now I did another nice experiment with halogenate ions. I recently purchased a kilo of NaI (sodium iodide). I did an electrolysis experiment with this, in an attempt to make iodate. This also works well, although it is a little bit more messy than making KBrO3. Dissolve NaI in water and electrolyse with a graphite anode and a metal cathode (but graphite is even better in order to prevent metal contamination). At the anode, a heavy brown compound is formed. This is iodine, which dissolves in the iodide solution. Also some solid iodine is formed (turbidity), but this quickly dissolves on shaking. When this process is allowed to run for a while, then around the anode, the liquid becomes really dark brown and heavy stuff sinks to the bottom. With shaking, however, the liquid quickly becomes light yellow again. No oxygen is formed at the anode. At the cathode, hydrogen gas is formed. The voltage, which needs to be applied is quite low (somewhere between 4 and 5 volts). When too high a voltage is applied, then oxygen gas is formed at the anode besides iodine, and the anode pulverizes in that case. When the electrolysis is running for a very long time, then oxygen is formed anyway at the anode, indicating that the concentration of iodide is dropping and that quite some iodate is in solution already. When this point is reached, you can stop the electrolysis, but then you loose quite some unreacted iodide. So, you need to have a compromise between iodide lost and pulverization of the anode. This definitely happens, when a lot of oxygen is formed at the anode. What happens is the following: cathode: 2H2O + 2e ---> H2 + 2OH(-) Anode: 2I(-) ---> I2 + 2e By mixing: I2 + 2OH(-) ---> IO(-) + I(-) + H2O 3IO(-) ---> IO3(-) + 2I(-) The iodide formed is electrolysed again. When I stopped electrolysis, I had a yellow solution. I added a very small amount of NaOH and then heated the solution. This assures that all hypoiodate and iodine is converted to iodate and on continued heating you also get rid of some water. To this, I added a hot concentrated solution of KCl. On cooling down this gives crystals of KIO3. KIO3 is only sparingly soluble in cold water. So, the method of making KBrO3 can also be used for making KIO3. KIO3 is not very energetic from the point of view of pyrotechnics, it is totally different from KBrO3 in this. KIO3, however, can be used for many other nice experiments and this electrolysis experiment can easily yield you several grams of KIO3 from a few grams of sodium iodide. ----------------------------------------------------- Now I have a question about the electrolysis with graphite. What causes the pulverization of the anode? I get this pulverization, as soon as oxygen is formed. Is this due to the bubbling? But if this is the case, then why don't I have this kind of problems with the cathode? All the time I had a lot of bubbling at the cathode and that is not pulverized at all. I have observed the pulverization of graphite anodes very often in all kinds of electrolysis processes, and this always has wondered me. Has anyone an idea?

Are you sure about the lead iodide. I do not think so. Lead iodide is soluble fairly well in warm water. One dust-wipe with a warm damp tissue would dissolve an appreciable amount of the (yellow) paint. I hardly can image that lead is used in modern house-paints. It still is used in artists-paints, but that is not used for painting wood parts of buildings. But for the OP, if you really worry, have lead levels checked in your blood, but I personally would not worry too much. But for the future I would be very careful not to use painted wood or at least have a good check on its source.

To my opinion the best thing you can do is respond and tell that natural vitamin C and man-made vitamin C are the same and that you cannot make any difference between them. If there is a difference, then it wil be due to impurities. But good refining methods certainly can make the natural product as pure as the synthetic product. Ascorbic acid crystallizes well from water, so it should not be that hard to make a pure compound from natural sources.

Ferrous chloride dissolves approximately at 10 gram per 100 ml of water, according to my data. This compound crystallizes as FeCl2.4H2O and in solution it forms ions Fe(2+) and Cl(-). 10 grams of this compound is approximately 0.05 mol. This means that the solubility is approximately 0.05 mol per 100 ml, hence 0.5 mol/liter. For soluble compounds, there is the concept of solubility product. For FeCl2 this is (taking the value of 0.5 mol/l as the maximum amount which can be dissolved): Ksp = [Fe(2+)]*[Cl(-)]*[Cl(-)] = 0.5*1.0*1.0 = 0.5 mol3/l3 Remember, for [Cl(-)] the double value must be plugged into this formula, because we have 2 mol of Cl(-) ions for each mol of FeCl2.4H2O Now, the common ion effect comes into play. In HCl, there are LOTS of chloride ions. Suppose you start off with 30% HCl and half of its chloride remains, the rest is used up to form FeCl4(-). Then still 15% HCl remains. Taking into account the higher density of HCl-solution, 15% HCl roughly contains 170 grams of HCl per liter, which means roughly 5 mol/l. Plugging into the formula for Ksp: [Fe(2+)]*5*5 = 0.5 ==> [Fe(2+)] = 0.02 mol/l Using the molecular weight for FeCl2.4H2O and the fact that one Fe(2+) ion results in one FeCl2.4H20, this means that only approximately 4 gram of this can be dissolved in one liter of 15% HCl. I hope you grasp the concept of this computation. All the numbers I have taken are VERY rough estimates, but using this kind of reasoning you certainly can explain why you see solid FeCl2.4H2O. Due to the rough estimates I may be off by a factor of 2, but still then, the reasoning holds. So, the answer is 'yes', you can certainly have solid FeCl2.4H2O in your liquid. What I think is more important, however, here, is that you understand this way of reasoning. ------------------------------------------------------ Just for education, a related little experiment you can do without problem (you have the reagents, I'm quite sure ). Take 2 ml of water and dissolve some table salt in it. See how much you can dissolve. Next, take 2 ml of concentrated hydrochloric acid and see how much table salt you can dissolve in that. Try to explain the observation.

Aluminium metal actually is useless as electrode material. The oxide layer prohibits acceptable currents. Your potential may be high, but you have a huge internal resistance in your 'cell'. As soon as you connect a load to the cell, the voltage drops, due to the high internal resistance. Oxidizing solutions do not help you with this. Strongly acidic solutions (like HCl) may help somewhat, but these quickly dissolve the aluminium after an initial transient. I would suggest you to take zinc metal instead.

If you really are concerned about lead dust in the house, then clean it thoroughly. Because the lead compounds are not volatile (as opposed to mercury metal), a single good cleanup of the house by means of removing dust with a damp towel should be OK. Any carpets should be vacuumcleaned well. Of course, I cannot say much about past contamination, but also for that, I would not worry too much. If wood is burnt with some lead-containing paint on it, then the lead will be converted to lead (II) oxide, PbO. This compound, especially when calcined by strong heat, is completely insoluble. Getting this in your body mostly will result in unchanged excretion, but tiny amounts of course may be absorbed. If i were you, I would thoroughly clean the house in the way mentioned above, and with future wood to be burnt, I would not use any painted wood anymore.

Tritium indeed is an isotope of hydrogen with two neutrons, written as [math]^3H[/math]. The most powerful chemical explosive I do not know. As soon as I mention one, a slightly more powerful one can be mentioned. This is not such a relevant question. Whether the VOD equals 8000 m/s or 8500 m/s or whether 3000 KJ/kilo or 3500 KJ/kilo is released is not that interesting. What counts for explosives is their usefulness and the cost. E.g. TATP is totally useless crap, due to its sensitivity. Other reasons may make an explosive less useful, such as being very hygroscopic, being extremely toxic, having short shelf-life. Cost also is an important factor. Fulminating gold for instance is very expensive.

How do you perform your reaction? Just with the vessel in direct contact with air. I did my experiments in a cork-stoppered test tube and then air only enters the test tube very slowly. I'm afraid it is the same trouble as with CuCl and aqueous copper (I) compounds. FeCl4(-) is very stable, even at low concentration of chloride ion, so the iron (III) species is strongly favoured under your conditions. The situation is not as bad as with CuCl and the colorless CuCl2(-) solution, but still, it is annoying. So, if you dissolve your iron in a solution with iron (III) and HCl and you close the vessel, and you use excess iron, then you can get iron (II) in solution and keep it like this. But every time, when you open up the vessel to take out some, it will become colored yellow again. Even small amounts of FeCl4(-) make the liquid appear yellow quickly. You could delay the process by taking a can of 3M dust remover. This is a pressurized relatively inert gas (butane, dimethylether, or 1,1-difluoroethane), which is used for blowing dust from camera-lenses, computer keyboards, audiodevices, etc. Each time, before closing the vessel with the ferrous chloride solution you could blow in some of this gas and then very quickly capping the vessel. This strongly reduces the oxidation, but does not eliminate it. I only know one iron (II) salt, which stores reasonably well and that is Mohr's salt, a double-salt, FeSO4.(NH4)2SO4.6H2O, ferrous ammonium sulfate. This salt is not oxidized by air. Plain ferrous sulfate, FeSO4.7H2O in due time becomes covered by a brown crust of basic ferric sulfate. An unoxidized ferrous salt is very pale green/blue, see http://woelen.scheikunde.net/science/chem/compounds/ferrous_amm_sulfate.html

In fact, all chemical reactions depend on temperature. This also is true for nuclear reactions, but for them, temperature sensitivity is on a totally different temperature scale. At ordinary temperatures (up to thousands of degrees), the nuclei are shielded effectively and then nuclear reactions are not affected. But at tens of millions or hundreds of millions of degrees there certainly is a dependence on temperature, just as there is in the range of hundreds of Kelvin for ordinary chemical reactions.

No, at least not in appreciable amounts. I oce read something about volcanos, which could give off trace amounts of elemental chlorine (which is converted to chloride quickly anyway), but from a practical point of view, I would so 'no'. Correct.

This link may give you some answers: http://chemlab.truman.edu/CHEM121Labs/Electronegativity.htm Electronegativity is related to bond energies. Bond energies can be measured indirectly by measuring net energy production or absorption by formation or breaking of certain bonds (although it is not easy at all to do it accurately).

The water plays an essential role. HIO4 is metaperiodic acid. With water, orthoperiodic acid is formed H2O + HIO4 <---> H3IO5 I'm not sure, it may even go further to H5IO6. Now, you have multiple HO-.. groups per molecule of acid, which can form a bridging (cyclic) periodate ester with more than one ester-group.

Bluenoise has a very good point here. I think it is VERY unlikely that initially there was elemental chlorine. Elemental chlorine is way too reactive to survive in nature for more than a few hours. So, even in very early times, I only can imagine the presence of chloride, not of chlorine. So, the answer to the OP is 'no'.

Polygon? Or do you mean univariate polynomial y = P(x)? It simply boils down to solving a non-homogeneous set of linear equations Ac=b. Here the vector c is the vector of polynomial coefficients, the matrix A is built up of all powers of x and the vector b is built up of all values of y. Keep in mind though, that in general the matrix A can become very ill-conditioned for a large real set of points i from 1 to N, {(xi, yi)}. In numerical solving this may give big troubles, even for moderate degrees of the polynomial (e.g. degree equal to 20 is at the border for standard IEEE 64 bit precision, C's double).