woelen

In order to solve this problem, a few questions arise 1) Is there any gravitational pull and hence, is there a gravitational acceleration towards to ground? What is the value of this pull (9.8 m/s2)? 2) The speed of light is considered infinite? 3) The size of the ball can be neglected, i.e. it can be considered a point mass? Given these assumptions, the problem is fairly simple. Compute where the ball is after one second: dv/dt = -9.8 m/s2 ==> v(t) = v(0) - 9.8t = 13 - 9.8t (m/s) dh/dt = v ==> h(t) = h(0) + v(0)*t - 4.9*t*t = 0 + 13*t - 4.9*t*t (m) Plug in t = 1. This gives you the height and the vertical velocity. Now derive an expression for the distance of the shadow from the pole as function of h. This involves goniometry and is left for the OP as an exercise. You get some expression as function of h, lets call this d(h). The velocity at which the shadow is moving along the floor can be written as dd/dt = d'(h) * dh/dt = d'(h)*v, where d'(h) is the derivative of the function d with respect to h. Plug in the value of h at d'(h) and take the value of v, both at t = 1 second. That gives you the result.

To my opinion you are doing foolish and k3wly things. You light a mix of almost 70 grams of chems, without the chance to run away? You don't use fuses or some other mechanism to delay the ignition? I don't want to be rude, but if I'm correct with this, then I think you are a fool! If you do take some safety measures, then please elaborate on them over here and then I'll take back my rude words.

Are you sure? I had a bottle of sulphuric acid with just a standard flimsy plastic screwcap. The bottle itself was made of glass. After a few months I noticed that the cap was totally eaten away, the acid had absorbed moisture from the air and the level of the liquid had risen considerably (it almost flew over the rim!). I would suggest to take a glass bottle and use a thick plastic cap. At the drugstore, where I buy my acids, both 30% HCl and 52% HNO3 are sold in soft plastic bottles with ordinary screwcaps, but the 96% H2SO4 is sold in glass bottles with a heavy duty several mm thick screwcap.

I agree with this. It is imaginable that a gaseous compound with an icreadibly large molecular weight has a larger density than the least dense liquids and solids. But from a practical point of view, I indeed do not know of any gaseous compound, with these properties.

I left my selenium in the open air in a warm place for a few days, taking care of that no dust can come in(remember, it was summer when I did my experiments) and it is nicely dry and has no nitrous vapors clinging on it. Nor does it have a bad smell. All nasty stuff has outgassed quite well. But, there is another bad thing with my sample. It has become darker . Not much, but definitely darker than when I made the picture for my web site. I'm afraid that in the long run, red selenium simply is not stable when it is dry and that if you really want it in your collection you have to replace it every year or so . If my sample indeed becomes dark, then I'll stop with this red selenium stuff and then I just store the grey corpuscles and keep my nice picture as an in memoriam for the red allotrope . I do not want to keep a wet sample under a dilute acidic sulfite solution.... and besides that, how long will that last ? But why don't you want to have some nitric acid around? If you store it in a well capped bottle and put the bottle in a plastic bag, what can happen with it? In my home lab I have a few liters of concentrated acids and a few liters of solvents around. Of course, you should not store these chems at places where you sleep, prepare your food or are living for several hours a day.

This problem cannot be solved with this information. Where does the oxygen remain? Is the oxygen removed as water (H2O). With these reactions, however, the oxygen also could be carried over as COCl2 (phosgene), which is a very common oxidation product when an oxygen-containing haydrocarbon is broken down in a stream of chlorine.

No, there are other stronger oxidizing agents, such as PtF6, which is capable of oxydizing oxygen. Another really strong oxidizer is sodium perxenate, Na4XeO6.8H2O. But indeed, ozone is a very strong oxidizer, similar in strength as peroxodisulfate (redox potential just over 2 volts). @budullewraagh: Permanganyl fluoride is not MnO4F, but MnO3F, with a direct Mn-F bond and Mn in its +7 oxidation state. I recently discovered that I actually made some of this stuff in one of my home experiments, it is a green volatile compound. Just google on permanganyl fluoride and you'll see a few links, where the formula MnO3F is given to this compound. I think that MnO4F does not exist. You can see this green stuff in the following link: http://woelen.scheikunde.net/science/chem/exps/KMnO4+NaF+H2SO4/index.html When I did the experiment I did not know it, but some literature study has shown me that the green compound is MnO3F.

Indeed there is a compound NO3F, but not a compound MnO4F. The highest oxidation state in which manganese exists is +7 and that is in MnO4(-), Mn2O7 and MnO3F. NO3F indeed is, as Jdurg mentioned, O2NOF, with F having oxidation state -1, the middle O having oxidation state +2 and the left two O's having oxidation state -2 and the N having oxidation state -3. This is a remarkable compound with oxygen strongly oxidized by the fluorine.

I agree with you. I also wish I had such a book. I have "Chemistry of the Elements" by Earnshaw and Greenwood. It is a nice book, but it covers the lanthanides and actinides only as a group. On one side, this book covers very rare and strange compounds, on the other hand, it does not even help me with the observations I do in many of the experiments I do with my very limited resources. Why did well-equipped labs not investigate such observations? That is what really irks me.

I'm quite sure your teacher is right in this case. There indeed is no known compound, capable of oxidizing fluorine. Fluorine only exists as element in oxidation state 0, or as fluoride in oxidation state -1.

No, you will not get CuO (or Cu(OH)2) in this way. You actually need the cathode material. At he anode you get Cu(2+) ions and with OH(-) ions, formed at the cathode, you get your Cu(OH)2. Anode: Cu --> Cu(2+) + 2e (absorbed by power supply) Cathode: 2H2O + 2e (given by power supply) --> H2 + 2OH(-) When the liquid is mixed (and it does on its own, due to the motion, induced by the rising gas bubbles), then the hydroxide and copper ions form a precipitate.

In your situation with lots of KNO3 at hand, forget about the Na2SO4. I mentioned this as just being a rather inert salt, which does not form complexes with copper (II) ions, nor with copper (I) ions. For most people, Na2SO4 is easier to obtain than KNO3, so I mentioned that. Separating CuO from the liquid is easy. CuO forms a precipitate and sinks to the bottom. Decant the solution, add a lot of water, let precipitate settle again and decant the water again. Doing this twice gives you wet CuO, which can be dried in a few days on a piece of filter paper, which you put in a warm and dry place. A very nice source of chemicals are pottery and ceramics suppliers. They have many metal oxides, metal carbonates and metal sulfates at reasonable purity and really cheap. If I were you, I would try to grab some metal salts from such a shop. With copper (I) I mean compounds of Cu(+), e.g. CuCl and Cu2O. With copper (II) I mean compounds of Cu(2+), e.g. CuCl2, CuSO4, and CuO.

Xeluc, finally I've done an experiment, which I promised you to do. I made some nice white CuCl by adding a concentrated solution of a copper (I) salt to a lot of water. For you the preparation of the snowy white stuff is quite familiar now I guess. Next, I decanted the very light blue solution above the white crystalline precipitate and I rinsed with water. Next, I decanted again, leaving a wet white solid with a very thin layer of water above it. This stuff, I heated until all water has evaporated. The results are really disappointing. I get dirty brown solid, with even black crap in it. Remarkably, on cooling down, the stuff became a little lighter and the black color disappeared. Finally, I had a light brown solid, but certainly not the beautiful white solid I had before. Now comes the most remarkable part. To this brown crap I added a little water again. And guess what happened . The solid becomes snowy white again within seconds. This REALLY surprised me. As expected, the solid does not dissolve. -------------------------------------------------------- Next, it became time to sit down and think . I have the following theory on these observations. 1) The nice white stuff is not as pure as we think it is. We always make it from a strongly acidic solution, which also contains quite some Cu(II) ions. 2) On dilution with water, the CuCl crystallizes and the solution becomes waekly acidic with low chloride concentration. The dark brown multivalency complex cannot exist under these conditions anymore, the copper (I) precipitates in CuCl, the small amount of copper (II) goes in solution as blue aqua ions. 3) Some of this blue copper (II) material is attached/embedded in the CuCl-crystals. This makes them optically very white. Some yellow/brown impurities are masked by the blue color. This principle of optical whiteness also is deployed by certain brands of clothes washing soap to make things look brighter and cleaner than they actually are. Our CuCl crystals may look so beautifully white, due to this copper (II) impurity. 4) Also some acid is embedded in the precipitated CuCl crystals. 5) On heating to dryness, the copper (II) in the mix forms brown copper (II) chloride. The black stuff may be transient copper (II) oxide, which disappears again because of some HCl still present in the crystal mass. I also go quite some white fumes during the heating. 6) What is left behind is copper (I) chloride, with some copper (II) chloride and possibly a very small amount of acid left. 7) On addition of water, the copper (II) chloride becomes very light blue again and partly goes into solution. With the traces of acid left, the liquid does not become turbid. The CuCl again looks snow-white again. I invite you to try this yourself. It is really remarkable to see all of this. If you heat the white crystals be careful to use heat resistant glassware. I used a wide testtube.