woelen

A very sensitive reaction, which only is positive for nickel and not for iron, copper, cobalt, chromium and most other metals is the following: Add an excess amount of a solution of NaOH to the liquid. This gives a precipitate. Then add a solution of a persulfate, such as Na2S2O8 or (NH4)2S2O8 or K2S2O8 (Feinätzkristalle) to the precipitate. If it contains even a small amount of nickel, then it turns black like carbon within a few seconds. These persulfates can be purchased at electronics parts stores as printed circuit board etchant. They are very nice oxidizers in aqueous solution, however, not suitable for pyrotechnics. This is a very nice result with the crude starting materials you have used. If the solution was very pale green, then there was no strong nickel contamination. Nickelic solutions have a much more intense green color.

This is funny. I also once did a similar experiment, but then with electrolysis. Take a 12 V voltage source (e.g. old PC power supply). Take a 10 Ohm power-resistor and place this in series with the voltage source. Assure that the resistor is not becoming too hot, otherwise use two parallel resistors of 22 Ohm. Perform electrolysis of a concentrated NaCl solution. In parallel to the NaCl solution, also take a 10 kOhm resistor in series with a 10 uF capacitor. The voltage over the 10 kOhm resistor will be 0 on average, but there is a considerable AC-component. If you connect this to the LINE input of an audio amplifier, then you'll hear the electrolysis proceeding. Quite impressive. It is a kind of roaring noise. Be careful here and use low volume and then turn up the volume. The signal level can be quite high and you may damage your loudspeakers if performed carelessly.

Oh yes, you can easily remove them without damaging your skin. You need to take very dilute HCl (let's say 3% by weight or so) and add a pinch of sodium sulfite or sodium metabisulfite to the acid. This gives the lovely smell of SO2. Apply this to the stain and almost immediately the stain is gone. Next, rinse with water. The precise amount of Na2SO3 does not matter. Na2SO3 is fairly benign and hardly harms the skin, not even at high concentration and exposure for minutes. An alternative is to prepare a 2:1 mix of 3% H2O2 and 10% HCl and apply this to the stain. Equally effective as the previous one and also harmless to the skin. Also this should be rinsed away after applying it. With the H2O2 mix you should be sure not to use higher concentration of H2O2. Only thing is that these dilute acid mixes should not be applied on damaged skin or wounds. That also is not that harmful, but the stinging is quite painful in that case.

Have a look at http://www.emovendo.net for pure elements. Iron is particularly cheap in the form of small scales at very high purity. These do not as quickly rust away as powder and with patience they can be dissolved in hydrochloric acid. You also can buy iron powder, but that should be stored in a VERY well closed container, otherwise it will soon degrade to rust. The powder, used for magnetic experiments is less suitable. It contains a lot of contaminants in order to make the powder less susceptible to corrosion. Very pure iron is not an easy element on storage, unless it is in the form of small scales. These scales look as follows: http://woelen.scheikunde.net/science/chem/compounds/iron.html I indeed have purchased quite some metals at high purity for the special purpose of chemistry experiments (aluminium, zinc, tin, lead, iron, nickel, chromium, silver). These metals are not that expensive. The only metal, which can be found in everyday items at high purity is copper as electricity wire. All other metal items simply have too much impurities. Besides the cheap metals mentioned above, I have purchased gallium, indium, antimony, vanadium, rhenium, ruthenium, molybdenum, niobium and cobalt for chemistry experiments. Sometimes it also is worthwhile to buy salts of metals. Ceramics/pottery shops have salts or oxides of many common metals at remarkably good purity (e.g. copper, cobalt, iron, nickel, chromium, vanadium) for just a few bucks.

Indeed, you have to be VERY close to equilibrium. But... this only is true for strong acids and strong bases. When weaker acids are used, then this need not be the case anymore. One can make so-called buffer solutions, which are designed to have a certain pH and to which quite some acid or base (beit strong or weak does not matter) can be added without a large change of the pH. Such buffer solutions also can be designed for pH close to 7.

The Al is oxidized. You start with Al-metal (oxidation state 0) and end up with Al2O3. Here, the metal has oxidation state +3 and the O has oxidation state -2. In the OH(-) ions, oxygen also has oxidation state -2. So, the Al is oxidized and not the OH(-).

I've been out of town during the Christmas days. I wish a very good New Year for all of you.

I've read all your observations and I also have a hard time to explain all these things. I really urge you to try to find sources of pure iron. These scratchpads and steelwool things and so on are so increadibly impure (at least where I live). They contain other metals (sometimes even small amounts of tin) and they contain a lot of grease-like stuff and detergents. Not good at all to do experiments with them. In acidic environments these detergents are converted to the corresponding organic insoluble acids (YT, remember what happens when acid is added to sodium benzoate, a similar things happens with the detergents in soaps). Other metallic impurities may give rise to all kinds of colors (dark green for chromium, green for nickel, white flocculent precipitate for tin). @jowrose: Take some of your green powder and add some of this to water and shake well. Do you get a clear solution? Add some HCl. Does this make the solution clear? If acidification results in a clear solution, then you don't have detergent impurities nor tin impurities. Also another word of warning: I received a bottle of 50 grams of reagent grade FeCl2 (anhydrous) from an old German lab. This bottle was still sealed, but it was almost 15 years old (it was prepared in the former GDR, just before "Die Wende", 1989). When I opened the bottle, the contents is totally brown/yellow and it does not dissolve clear anymore. It also does not reduce acidified dichromate anymore. So, all FeCl2 is oxidized in these 15 years, even in the sealed bottle . So, if even such a pure reagent grade compound is oxidized in a sealed bottle, then you can imagine how hard it will be to make pure FeCl2, which keeps for a reasonable time. The stuff I have can best be regarded as basic ferric chloride.

What the OP does has nothing to do with electrolysis. The carbon electrode becomes VERY hot and at the water/carbon interface there will be an hefty main reaction: C + H2O --> CO + H2 At the temperatures involved, CO2 is not stable and CO will be formed. This is a dangerous gas and may kill you without you even noticing anything (it is odorless, colorless and tasteless and does not give any irritation or discomfort while breathed, as opposed to most other toxic gases, like Cl2, HCl, etc.). So, indeed, here you cannot talk of anode/cathode, there simply is a rod of carbon at red heat in the water. Before the above mentioned reaction occurs, however, the temperature should be really high. I don't know the exact figures, but it is not a matter of 100 degrees centigrade or so. It must be 1000 or even 2000 degrees. The entire carbon electrode disappears. It, together with the water is converted to gaseous products only. You cannot give a precise relation between the amount of electric power or charge taken and the amount of hydrogen produced. With true electrolysis there is a very precise relation, because each molecule of hydrogen requires two electrons. Here, however, the driving force is heat, not charge/electrons. ANy heat source would do, you now just use the electric current for heating the carbon rod.

sodium hydroxide is corrosive: This is corrosive and fairly toxic as well. It destroys organic (including human) tissue, even at moderate concentration. This makes it toxic. hydrochloric acid is corrosive: It is very corrosive and very toxic, when concentrated. When diluted it still is quite corrosive, but its toxicity quickly lowers on dilution. The concentrated acid, however, gives an intensely choking fume. Never take a deep breathe near the open mouth of a bottle of concentrated hydrochloric acid. You'll regret such a thing . barium chloride is harmful maybe toxic: It is VERY toxic. nitric acid corrosive: VERY corrosive and intensely poisonous. It also is a strong oxidizer. silver nitrate is corrosive: corrosive and toxic, and a strong oxidizer. It carries the harmful cross, the corrosive sign and the sign for strong oxidizer. http://ptcl.chem.ox.ac.uk/MSDS/SO/sodium_hydroxide.html http://ptcl.chem.ox.ac.uk/MSDS/HY/hydrochloric_acid.html http://physchem.ox.ac.uk/MSDS/BA/barium_chloride_dihydrate.html http://ptcl.chem.ox.ac.uk/MSDS/NI/nitric_acid.html http://ptcl.chem.ox.ac.uk/MSDS/SI/silver_nitrate.html

All of you are right in some sense, but let's be practical and realistic. For the average home chemist it is VERY difficult to make sodium metal in a safe way and to isolate it. I myself do quite a lot of 'mad scientist' experiments, but the making of sodium at home I do not even consider. It simply is not worth the risk. Imagine what happens in case of an accident, with molten NaOH or NaCl sprayed around. Or with the amalgam method and trying to purify this. Think of the toxic mercury. Forget about making sodium metal at home. If you can't find a source of it, accept that, or do a better job finding one. There are many other fun experiments to be done at home as well.

It indeed has to do with the logarithmic scale of the pH. I'll clarify with an example. Suppose you have one liter of liquid, with 0.1 mol of acid (e.g HCl) in it. Now you add drops of concentrated base to this liquid. A concentrated solution of NaOH contains approximately 10 mol/l of NaOH. Each big drop of 0.1 ml contains 0.001 mol of NaOH. Now suppose you add these drops to the liquid and after each drop you mix well. What you get is the following: After 1 drop 0.099 mol of acid remains. After 2 drops 0.098 mol of acid remains. After 3 drops 0.097 mol of acid remains. ... After 98 drops 0.002 mol of acid remains. After 99 drops 0.001 mol of acid remains. After 100 drops the liquid is neutral After 101 drops 0.001 mol of base is in the liquid After 102 drops 0.002 mol of base is in the liquid ... After 200 drops 0.100 mol of base is in the liquid. Now take the -log10() of the concentration of H(+) ions and you see what I mean. For simplicity neglect the increase in volume of the total liquid. It only increases from 1.00 to 1.02 liters, so that is not much of influence. So, pH as function of number of drops (at 3-digit accuracy): 0 --> 1.000 1 --> 1.004 2 --> 1.009 3 --> 1.013 ... 10 --> 1.046 ... 20 --> 1.097 ... 50 --> 1.301 ... 80 --> 1.699 ... 90 --> 2.000 91 --> 2.046 92 --> 2.097 93 --> 2.155 94 --> 2.222 95 --> 2.301 96 --> 2.398 97 --> 2.523 98 --> 2.699 99 --> 3.000 100 --> 7.000 101 --> 11.000 102 --> 11.301 103 --> 11.477 ... 110 --> 12.000 ... 200 --> 13.000 You see that the pH jumps from 3 to 11 with just two drops, while the first few drops and the last few drops hardly have any effect on the pH. This is just a property of the logarithmic scale. You add base linearly. ------------------------------------------------------------------------------- If you did the same experiment with a weak acid, then the curve would be less steep around the neutral-point. This is because for weak acids you have an equilibrium shift when base is added and then not all base will be used for neutralization of the acid. The curve becomes 'softer/smoother'.

To my opinion it is plain stupid to play with your health like this. Here we are talking about short-term effects, but what are the long-term effects of repeated exposure?

I would suggest you keep the niobium in its nice container. Since three weeks Emovendo sells 30 grams of 99.8% niobium for $22 in the form of such sheets, as I show on my site. This is my source of the niobium. For my element collection I have a nice rod of the metal, but for chemistry experiments the sheets from Emovendo are very nice and affordable (although not really cheap).

Xeluc, thanks for this information. This was something I did not know, so I learned a lot about niobium chemistry. I updated my webpage and added the information from your link to my page.

I did an electrolysis experiment with 99.8% pure niobium metal as anode. I tried this, in an attempt to dissolve the very inert metal. My reasoning was that the metal either dissolves or allows all kinds of gases (such as chlorine) to be produced at the anode. The latter would be very nice, having a metal anode, which does not pulverize with electrolysis. In reality, however, something really stunning happens. This is so weird that I decided to devote a web page to it. The experimental results can be read here: http://woelen.scheikunde.net/science/chem/exps/niobium/index.html

I also tried this. It simply quenches, no spectacular reaction at all. It is not that the ribbon sucks, but the piece of metal, relative to the amount of water is too small. It probably would work if you had a larger lump of very hot and burning magnesium and put that in a thin layer of water, such that it cannot be totally covered with water.

Don't try to do that with ammonium dichromate. You need K2Cr2O7 or Na2Cr2O7.2H2O. In fact, ammonium dichromate is not a very interesting chemical, except for the volcano demo and some pyrotechnics things. For all aqueous and solution chemistry, the sodium- and potassium salt are MUCH more useful. The ammonium ion is prone to oxidation, when mixed with CrO3 and that can give rise to really dangerous situations. Btw, why do you want CrO3? Almost all experiments with CrO3 can also be done by dissolving K2Cr2O7 in dilute H2SO4. CrO3 is a pain on storage. I have 25 grams of this, but within a year it has become a bad black/brown, sticky and exceedingly corrosive mass. My K2Cr2O7 already is 20 years old and it still is a beautiful bright orange crystalline solid, which shows no deterioration at all. My (NH4)2Cr2O7 also keeps very well. Even, although I have CrO3, I hardly use it, because K2Cr2O7 is a much safer and easier to handle alternative. CrO3 is so corrosive, that it ignites organics and immediately destroys your skin on contact. Beware of all hexavalent chromium compounds. They most likely are carcinogens.

What is sds? Explaining that may help others in solving your problem

You are talking about acidified solutions of dichromate. Then you have the following unbalanced equation: The dichromate requires acid to be present in order to oxidize, so you'll have H(+) at the left. The dichromate is an oxidizer, so the half reaction has electrons at the left. The reaction products are water and Cr(3+) ions. Cr2O7(2-) + H(+) + e --> Cr(3+) + H2O Now, how do we come to the precise reaction equation? It is just a matter of good bookkeeping. At the right, we have Cr(3+) ions. At the left we have two Cr-atoms in the dichromate ion, at the right we need two Cr(3+) ions. Now, we are left with the oxygens from the Cr2O7(2-). There are seven of them, so we need 14 H(+)'s in order to compensate for all of them. The result is 7H2O at the right. Now, the equation is balanced for Cr, H and O. We still have to balance for charge. At the left, we have a charge, equal to +12 for the Cr2O7(2-) ion and the 14 H(+) ions. At the right we have a charge, equal to +6, due to the two Cr(3+) ions. At both sides we need the same charge, so we need to add 6 electrons at the left (charge of -6) in order to make the equation correct for charge as well. Try to do this as an exercise for the reaction of pervanadyl ion to vanadium (III) ion. If you can derive the correct half-equation for this, then you understand the concept. VO2(+) + H(+) + e --> V(3+) + H2O Now, the question, why a negatively charged ion can change to positively charged ions due to reduction. This answer is not that difficult. The charge of the ions involved does not really say anything about the possibility of reaction. Look at the left-hand side of the equation. A lot of H(+) ions go in as well. The electrons are not really present as free entities, they are just introduced for bookkeeping purposes. In reality, they are passed immediately from reductor to oxidizer and there are multiple mechanisms. Google for "outer-sphere" and "inner-sphere" redox reaction mechanism. With these concepts, a lot more will be explained on the mechanisms and I think things will become more clear to you. Here follows a final example, which may be interesting to try. It also goes from the yellow positive VO2(+) ion to the bright blue positive VO(2+) ion: VO2(+) + H(+) + e --> VO(2+) + H2O Try to balance this one. If you can balance this correctly, you really understand this . I've done this reaction many times and it is very neat to watch it, going from deep yellow through all kinds of green until a final bright blue color is obtained.

There are a number of nitrogen/oxygen groups in organics: nitrate = -O-NO2 nitrite = -O-NO nitro = -NO2 E.g. CH3-ONO is a gas, called methylnitrite CH3NO2 is a colorless oily liquid, called nitromethane. The compound CH3-O-NO2 is called methylnitrate. This is a VERY unstable compound, really hard to synthesise (if at all possible).

It cannot be made easily from household materials. What do you want to do with it? If you want more info, you may send a PM to me. I do not publish sources for the chem over here, k3wls are reading this also unfortunately. PM me with what you want and I'll see if I can help you.

Forget about making sodium in your own house. You are asking for deep troubles and accidents. It simply is too dangerous, unless you have some really good equipment and a LOT of experience with experimenting. I myself am doing quite some chemistry experiments at home, but extracting sodium from NaOH and/or NaCl I do not attempt. I appreciate my health and my life too much to spoil it with such an experiment .

In fact, all noble gases form "compounds", which even have a distinct formula. A well known "compound" is C6H4(OH)2.Ar This is a so-called clathrate. C6H4(OH)2 is hydroquinone (p-dihydroxybenzene). When a solution of hydroquinone in water is allowed to evapotate slowly in an atmosphere of argon under high pressure, then argon atoms become trapped in holes of the crystal lattice of hydroquinone. So, this is not a chemical compound, but a physical compound. The holes in the crystal lattice of hydroquinone are sufficiently large to contain a single atom of argon. If such crystals are added to water, then they dissolve and bubbles of argon escape again. Clathrate compounds are quite common. Methane molecules can be trapped in a crystal lattice of water (ice) and this "hydrated" methane exists in large quantities at the bottom of the oceans. Yet another example is "hydrated" chlorine, Cl2.8H2O. This is chlorine, trapped in a crystal lattice of water (ice) and it can be made by freezing water in a chlorine atmosphere at 3 atmosphere pressure at a temperature of a few degrees below zero C. Even more interesting clathrate compounds (a.k.a. cage-compounds) are derived from fullerenes, C60. Such a fullerene molecule, which has the shape of a soccer-ball is hollow. Inside, a small molecule or atom can be trapped and indeed, such compounds have been made, e.g. C60.Au.