woelen

Yes, I have the same skepsis as you have. I have quite some experience with electrolysis experiments and having a few amps of current already is a lot. With 15 amps of current, the cell would be red-hot, due to resistive effects (dissipation) and the anode of the electrolysis cell would have to be replaced every few minutes, unless it is made of a very inert platinum-tantalum-iridium alloy, but I doubt that such an expensive anode would be used. @Jdurg: I think that most of the increase of weight is due to water, sticking to the piece of palladium. I certainly believe that some hydrogen is absorbed and even if 0.1 mole is absorbed and you heat the sample of palladium under water, then you can see a LOT of bubbling already.

It might be that all this reasoning about shells want to be half full and so on is correct, but there are counter examples. In the lanthanides there is even more complicated behavior with oddities in between (the f orbitals come in play then). So, I agree with the answers, because they make it easy to remember which are the oddities for the transition metals. But now look at the second row of transition metals (elements Y ... Cd), or look at Pt and Au, over there this reasoning does not work anymore. There only is one way to really understand all these oddities and that is numerically solving the quantum mechanical wave equations for all electrons in an atom, also taking into account relativistic effects. This is a hell of a job and requires a lot of computational power, but it has been done for us already. Only then you see the very subtle differences in energy levels. Probably this is not a nice answer in terms of easy understanding, but sometimes you have to take things as they are. Simple underlying physical laws can give rise to really complex behavior, which only can be understood by performing massive computations. Another nice example of a very simple system, following very simple laws is a three-body mechanical system with Newtonian laws of gravity. Such a simple system exhibits amazingly complex behavior, which only can be understood by simply running the simulation for a long time.

It is a matter ok knowing. You can assume that all ammonium compounds and alkali metal compounds are purely ionic, the compounds of the heavier earth alkali metals (Ca and further down) also can be regarded ionic. For the transition metals and the other metals, it depends on the counter atom/ion. Hydrated chlorates, perchlorates, nitrates and sulfates also may be regarded ionic in all cases. The anhydrous nitrates (e.g. Cu(NO3)2) can be covalent. Another rule of thumb is that the higher the oxidation state of a central element, the more covalent its compounds: E.g. CrO is purely ionic. Cr2O3 is intermediate, although it still can be considered ionic, CrO2 also is intermediate, and not purely ionic anymore. CrO3 is purely covalent. A similar thing holds for Mn in all its oxidation states. Also keep in mind, that ions themselves mostly are covalent entities. E.g. for KNO3, the salt is ionic, K(+) and NO3(-), but the ion NO3(-) itself is purely covalent, there is no nitrogen ion and oxide ion in this.

I was really startled to see what happened with the liquid in a period of 2 days. A beautiful blue/green crystalline mass had settled at the bottom of the test tube. The liquid above it was green with a yellowish hue. Apparently, still some air was trapped in the test tube and some oxygen made the liquid yellowish. I decanted the liquid and quickly rinsed the nice crystalline mass with some acetone. This preserved the crystal, but the color of the crystal quickly changed from the nice blue/green color to a purer green color, with a yellowish/brown hue . This is due to aerial oxidation, Fe(2+) being converted to Fe(3+). I made some pictures of the crystalline mass. It looks really neat. The diameter of the piece is approximately 1 cm. You can still nicely see the round bottom part of the crystalline mass, due to the round bottom of the test tube. In the first picture I also have drawn a small rectangle in the picture. That is close to the original color of the crystalline mass, before I took it out of the air-tight test tube.

You mean that you want to know a cold liquid, that remains cold? If that is what you mean, then I have to disappoint you. Such a liquid does not exist. Any cold object, exposed to room temperature, will finally reach that temperature. With good isolation you can make that process slow, but you cannot stop that. E.g. liquid nitrogen either becomes warmer (and then pressure builds up), or it evaporates and when all has evaporated, no liquid is left anymore.

Indeed, a LOT of hydrogen. This means 3 moles of electrons, or almost 300000 Coulombs of charge. Five hours is 18000 seconds, for simplicity let's make it 20000 seconds. Did you really use 15A of current for your electrolysis? That seems a LOT to me. Probably you must have had much more current, because not all hydrogen will be taken up by the palladium. The figures seem a little unbelievable to me.

No cats, but young children also sometimes can give rise to certain smells .

Yes, I know this. I also have this at my HNO3 bottles. I even had white needle-like crystals around the cap and the top of the bottle. I have scraped off some of the white stuff from a bottle, which had a lot of that stuff. I added this to a solution of NaOH and then I noticed a fairly strong smell of ammonia. So, I'm inclined to think that the white 'frost' is the ammonium salt of the acid (i.e. NH4Cl or NH4NO3). In the air, there always is some ammonia and this forms a solid with HCl-vapour and also with HNO3-vapor. If you have a bottle with quite some of the white stuff, just try it yourself and smell the ammonia.

Probably the best way is to precipitate the iron (II) sulfate in water and precipitate this with sodium hydroxide. You also add some H2O2 to oxidize all to Fe2O3. Next, filter the precipitate, rinse with some water and then dissolve it in HCl. Do not dry the precipitate. Next, precipitate another time with a solution of NaOH. Again, filter and rinse with water. Then dissolve again in HCl. That should give you pure FeCl3, dissolved in excess HCl. The reason for the extra dissolve/precipitate step is that the initial precipitate still contains quite some sulfate, due to co-precipitation. Making dry solid FeCl3 is not possible from aqueous solution. When the solution is boiled to dryness, then first the hydrate FeCl3.6H2O is formed. On further heating, this compound gives off water and HCl and a basic chloride remains. "Oxidizing flame" denotes a flame, in which the gas-mix is oxidizing (e.g. excess oxygen).

Suppose you have a two-compound reaction: A + B --> products A reaction only occurs if a molecule of A and a molecule of B meet at the same place. Now it is a matter of statistics. Because we have tremendous amounts of molecules of A and B, one can use with good approximation that the chance p(A) that a molecule of type A is at a certain place is proportional to [A], hence p(A)=k1*[A]. A similar thing holds for B, p(B)=k2*. Under the assumption that motion of molecules of type A and type B is not influenced by other molecules in the mix, one can state that the chance that both a molecule A and a molecule B are at the same place equals p(A)*p(B), hence it can be written as k1*k2*[A]* Now suppose we have a reaction of the type 2A + B --> products, and the reaction is a true 3-molecule type, which only occurs if 2 molecules of type A and one molecule of type B are at the same place. Again under the assumption of statistical independence, the chance that 2 molecules of type A are at the same place is p(A)*p(A), the chance that all three molecules are at the same place equals p(A)*p(A)*p(B). You see, now you get a power of [A], in my example this is power 2. With equal reasoning you can see that in general you get powers of the concentrations. ---------------------------------------------------------------------- Now suppose we take the reaction A + B --> products. This of course also can be written as 2A + 2B ---> 2*products. If you apply the same reasoning, then you would get another result. And here comes the great pitfall. In general, the stoichiometric reaction equations do not predict the rate of reaction, you really have to understand the reaction mechanism and you have to measure in a real practical setup the order of a reaction. Let's take my example 2A + B --> products. Suppose this reaction proceeds as follows: step 1) A + B --> AB (e.g. formation of a complex) step 2) AB + A --> products Step 1 has a rate d[AB]/dt = k1*[A]* Step 2 has a rate d[products]/dt = k2*[AB]*[A] Now you see that the real reaction dynamics require solving a set of non-linear differential equations. Now suppose that k1 is very large. This means that step 1 appears to run instantaneously. Suppose that k2 is not that large. Also suppose that [A] is in large excess to . What happens here is that on mixing A and B you at once get AB and the large remaining amount of A. If [A] >> , then the reaction still appears to be first order in its dynamics. The speed of reaction seems to be proportional to [A]*. If [A] is just a little larger than things become even more complicated. In general, reactions, where more than two molecules need to meet each other, are quite rare already. Reactions, where more than four molecules need to meet each other, I do not know an example of that. Almost all reaction take intermediate steps, such as in my example with 2A+B.

Don't trust the safety info from MSDS's too much. This info usually is strongly exaggerated. I only use them to get an idea about a compound's reactivity and to obtain general info like solubility, appearance etc. The safety info must be taken not too literally. Too my opinion magnesium nitrate is not very reactive. It is more corrosive than KNO3, however, because it is very hygroscopic and the wet salt corrodes things more than dry powder.

With the CuCl2 crystals, do not let all liquid evaporate on your nice and clean crystals. Crap and impurities build up in the liquor above the crystals, so when you let is all evaporate, your final part will be quite impure and it would be sad if that impure stuff mixes up with your pure crystals. So, I would advice to decant the liquid from your nice crystals and to have your nice crystals drying (that may be quite hard actually, CuCl2.2H2O is quite hygroscopic). The liquid can be put aside for other crystals, but these will be less pure and should not be mixed with your purer crystals. Drying CuCl2.2H2O can be done by putting your crystals in a watch glass or something like that, putting drying agent (anhydrous CaCl2, available at hardware stores) besides your crystals in another watch glass and put both these watch glasses in a closed vessel (a well-closed tupperware box is suitable very well for this purpose). The CaCl2 will dehydrate air in the box and due to the dry air, the CuCl2.2H2O will get dry, without loosing its water of crystallization. Within a few days you should get reasonably dry crystals. Have a look at this MSDS: http://ptcl.chem.ox.ac.uk/MSDS/IR/iron_II_chloride_hydrate.html Ferrous chloride is light blue/green, not white. If you want to be sure that your white crystals are ferrous chloride, then you could decant the liquid above them and add the crystal mass to some acetone, shake a little and then decant the acetone. Repeat this another time. In this way, you rinse off the water quickly and you make the crystals dry. Acetone is very volatile and your crystal mass is dry in no-time when you put the solid mass, wet from acetone, on a central heater radiator. Rinsing with clean water is not wise to do, because a lot of your crystalline mass may dissolve again. Before you put all your crystals in acetone, first try with a small amount, just to be sure that this does not spoil things. Btw, you should not use the acetone trick with your CuCl2 crystals. CuCl2 dissolves in acetone and forms a yellow/brown solution with a greenish hue! it is hard to say what all that added crap is. Indeed, usually a lot of other metals and sometimes carbon and even tellurium are added to the metal. If your precipitate is really crystalline (hard pieces with shiny edges), then it is decently pure, certainly enough for the kind of experiments you as a home chemist intend to do with it. If the precipitate is more like a paste and more flocculent, then I doubt whether it is FeCl2 and then I think it is sheer crap. In general, flocculent precipitates are VERY impure, due to a process called co-precipitation. Just as an example, suppose you have a solution of CuSO4 in water (light blue) and you add solution of NaOH, then you get a blue flocculent precipitate. On high-school one learns that this precipitate is Cu(OH)2. In reality, things are MUCH more complex. The precipitate also will contain SO4(2-) ions, Na(+) ions and who knows what more (oxo-species, water molecules). The physical processes in which such flocculent precipitates are formed, almost certainly cause other ions to be trapped (clathrated) into networks of -Cu-(μ-OH)-Cu- bridges. When a crystalline precipitate is formed (usually these are formed slowly), then the processes are quite different. Slowly, ions (or molecules) are deposited on the crystal and the deposition is such that the structure is nicely regular. That is why recrystallization can be used to make purer compounds. Flocculent precipitates make things impure, crystalline precipitates make things purer.

If there is no answer from the OP, then I only have one conclusion: yet another troll.

Getting really pure iron at some local store is not that easy. The reason for that is that very pure iron hardly has any commercial application. It is not sufficiently hard and strong for construction purposes and it rusts away easily. So, commercial iron is alloyed with all kinds of elements in order to prevent rusting and to make it stronger (e.g. chromium, vanadium, nickel, carbon). Steel wool, as the name says, is steel, which contains a LOT of other elements besides the iron (mostly carbon, but also other materials). From online shops, you can obtain iron at very high purity easily, but that requires you to buy online and pay with PayPal or credit card. The price should not be prohibitive http://www.emovendo.net/store/customer/home.php?cat=253 Iron is present over there at just $6 for 30 grams of ultrapure metal (99.98+ % purity). From this company I have some of the cheaper ultrapure elements (Fe, Se, Ag, Pb, Al, Sb). Shipping is only $1.60 to me in the Netherlands, so for you it should be even cheaper. Sometimes, eBay also has pure iron. My powdered iron is from eBay, at a price of EUR 7 per 100 gram, but purity is only 99.9+ % instead of 99.98+ %. That 0.1% can make quite some difference, especially if the impurity is a strongly colored species. Fortunately for my iron this is not the case, the impurity is mostly carbon, making the iron a little less reactive and causing the final pieces of metal not to dissolve as my picture shows (I assume they have a high carbon content). If you want to do nice metal experiments, then I also would suggest you to do experiments with chromium. That element gives beautiful blue and green solutions. Pure chromium also can be had from eBay sometimes. Emovendo, unfortunately is quite expensive on chromium, but with some luck, you can find cheaper sources of very pure chromium. Can you make a picture of your nice blue CuCl2.2H2O crystals? It is interesting to read that you get very pure CuCl2, which is not green. Indeed, slow evaporation of solutions is a good way to get nice crystals, the only disadvantage being that it may take VERY long. Most people are not that patient . The cubic crystals you get most likely are NaCl (from NaOH + HCl). Needle-like crystals? I cannot answer that, because I do not know the impurities from your iron and solutions.

You made me curious with all your iron experiments and I did an experiment with reagent grade HCl, 30% by weight and 99.9+ % pure iron powder, average particle size appr. 100 μm. I did a small spatula full of iron powder in a test tube and added some of the HCl. As soon as I did this, there was some slow bubbling, but definitely not vigorous. The liquid almost at once turned light yellow. On standing for a while, the yellow color turned somewhat more intense. Having the mix stand like this would take ages before all of the metal is dissolved. Next, I heated the test tube to appr. 90 C (not boiling, well below that point). At this temperature, the bubbling became much faster, but still not very violent. While the liquid was bubbling nicely, I capped the test tube loosely, in order to allow the gas to escape, but not letting air in. I heated for several minutes and it took me more than 10 minutes to dissolve the small amount of iron in the acid. At the point, where it was almost gone, I heated more strongly, until the liquid boiled vigorously and then at once, I took it out of the flame and capped it tightly (it still was loosely capped). In this way, I assured that no air is left in the test tube and on cooling down there is under-pressure. The liquid now almost is very pale green/blue, with a fairly large percentage of iron dissolved. All iron is dissolved as iron (II) species. On cooling down, the liquid remains clear. No flocculent precipitate can be observed. My observations can be explained as follows: First, in the HCl some oxygen is dissolved. This causes formation of the yellow FeCl4(-) complex (iron in oxidation state +3). When more iron dissolves and there is no air allowed to enter the test tube, air being displaced by hydrogen gas, then the iron (III) species is reduced to iron (II) species again. On longer standing, nothing changes anymore. So, the capped test tube contains a light green/blue liquid, which is totally clear. Some small pieces of metal remain undissolved. At a certain point the activity of the liquid (the acid) becomes so low, that the metal does not dissolve anymore. I attached a picture of my final result. It nicely shows the green solution. I could not reproduce your results of dark green liquids, flocculent precipitates etc. So, I'm afraid that your source of iron is quite impure. If it contains chromium, then a green color is understandable.

Jdurg and YT, this is very true for you, but I'm afraid vrus is in a country, where things are totally different. It really is remarkable, how difficult it can be to find even the most basic chemicals in certain parts of the world.

Aluminium can be used for electrolysis, but it is not very useful as an electrode for producing a current in a cell. What vrus attempts to do is not electrolysis, but make a cell, which produces some current. @vrus: Your electrolyte must have sufficient conductivity. Vinegar only is a weak acid and hardly conducts electricity. The internal resistance of your cell will become very large and make it useless. Can't you even get dilute HCl? If you can get that, then you can make a nice cell as follows: Dissolve as much as possible of zinc metal in HCl. As a substitute you may simply try a solution of NaCl instead of zinc, dissolved in HCl. That is even more simple and also will do. Dissolve as much as possible of copper metal in a HCl/H2O2 mix. Heat the mix for some time to get rid of excess HCl and H2O2. Put both solutions in separate beakers. Take a flexible tube, fill this with concentrated salt-water and close both ends with cotton-wool (which will be soaked with the salt-solutions). Immerse one end in the beaker with copper-solution and immerse the other end in the beaker with zinc solution. Dip a copper plate in the copper solution and dip a zinc plate in the zinc solution. Between the copper and zinc plates you will have a voltage with low internal resistance and that makes your cell nicely powerful. Try to go to a hardware store and see if you can find 10% HCl. I hardly can imagine that even such dilute acid cannot be obtained. A solution with 3 to 10% H2O2 should also be available. Easy to obtain chemicals might be copper sulfate and zinc oxide for you. They are available at pottery and ceramics stores. With these at hand your copper half-cell can simply be made by dissolving the copper-salt in water. The ZnO can easily be dissolved in dil. HCl. As an alternative for HCl you could try to find dilute H2SO4 as the acid in car batteries. With some luck you can obtain that at auto parts suppliers. Yet another nice alternative for HCl is NaHSO4, which is present in many toilet cleaners and also available as pH-minus for swimming pools. You should be capable to find one of these chems as your acid.

Ryan, I must correct you a little . Fluorine is not flammable. It supports combustion very well and many things can burn in an atmosphere of fluorine, but fluorine itself cannot burn in an atmosphere of oxygen. Caesium indeed is very flammable, but magnesium only somewhat. Of course it can burn with a bright and hot flame, but there are compounds, soooo much more flammable. Some examples: white phosphorous: This starts giving off smoke, as soon as it comes in contact with air. It quickly catches fire, even at room temp. With a warm hand you can ignite it. P2H4 (the phosphorous equivalent of hydrazine): This gaseous compound ignites as soon as it comes in contact with air, even in the cold. Most boranes also ignite on contact with air. For rocket related questions this may be interesting: http://www.ukrocketry.co.uk/forum/

Do they ship internationally ??

The problem is not clear to me. Which parameter is varied, what is kept constant?

Ay ay even experts do make blunders. In my previous post I mentioned that in 40 days, 1 mole can be made. One mole is just over 300 grams, so this makes 7.5 grams per day and not 0.75 grams per day . So, your yield is perfectly possible. The efficiency of your cell is not that good, but for home applications it is perfectly suitable and indeed your concept works nicely. The losses most notably are due to 1) Formation of oxygen at the anode 2) Back-reduction of chlorate and hypochlorite at the cathode to chloride. 3) Isolation of the chlorate from the brine with chloride salt