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If you have a saturated solution of Mg(OH)2, how much Mg(OH)2 will there be in solution, how many moles per liter? From this you can compute the concentration of OH(-) ions. Next, you can compute the concentration of H(+) ions, using the identity [H(+)][OH(-)]=10^(-14). That should give you the pH. Beware, in this computation I made a BIG simplification. It only works if [OH(-)] >> 10^(-7). For very low concentrations of OH(-) you need to solve a set of equations, resulting in a single quadratic equation. For your problem, however, the simplification is OK.

All of the earth alkaline metal chlorides are purely ionic, except the BeCl2, which is intermediate between ionic and covalent. When dissolved in water, expect the pH of solutions of all of them to be close to 7, except for BeCl2. This compound probably is somewhat acidic on dissolving in water, but I just reason from theory, I have no practical figures.

Using electrolysis, you can make reductors and oxidizers as strong as you like. Fluorine is the strongest chemical oxidizer known and cesium (probably) is the strongest chemical reductor known, but by means of electrolysis even stronger oxidizing and reducing conditions can be created. Redox reactions can be described by means of half-reactions, e.g. F2 + 2e ---> 2F(-) To such redox reactions, a potential is associated. For the reaction, mentioned above, a potential of approximately 3 volts needs to be created. If you electrolyse KF, then for the reduction of K(+) ions to K-metal also some potential is needed (of opposite sign). The total potential needed is somewhere around 5 or 6 volts (I'm too lazy too look up the potential for K(+) + e ---> K ). Usually some extra potential is needed, due to resistive losses and due to an effect called overpotential, which requires a few tenths of additional voltage for a reaction to proceed at an acceptable rate. As you can see, with a voltage of e.g. 10 Volts you have an increadibly strong reductor/oxidizer pair at the cathode and anode. Using such a pair, one can make cesium metal and fluorine gas (provided the correct equipment is used and the reaction products are not allowed to react with material in their neighbourhood). The strongest oxidizer, which I have at home and which is readily available is sodium persulfate, potassium persulfate and ammonium persulfate. The redox potential for the reactoin S2O8(2-) + 2e --> SO4(2-) equals 2.0 volts. This oxidizer is capable of oxidizing Ag(+) to Ag(3+), Mn(2+) to MnO4(2-) and Cr(3+) to CrO4(2-).

Yes, but with the copper/chloride combination there is something special. That mix is not very powerful, yet capable of dissolving aluminium. In fact, just plain table salt and some non-acidic copper (II) salt are much more effective than both HCl and NaOH, it is really remarkable. If you have some CuCl2 left from your copper experiments, then try the following: - Prepare a solution of table salt, quite concentrated. If it does not become nice and clear, then add a few drops of HCl, but not too much. If you add some aluminium household foil to this solution, then it does not react. Even with a small amount of HCl in it, it does not react. - Also dissolve some CuCl2 in that salt-solution. It need not be acidic, but if some acid is left in the CuCl2 it does not hurt. - Immerse some Al-foil in the copper(II)/salt solution. Be careful, this reaction is quite vigorous at once!! I made a web-page about this experiment: http://woelen.scheikunde.net/science/chem/exps/cu+al/index.html A really nice experiment is the last paragraph in smaller font. This, however, it quite dangerous, so be careful!!!!! You can do this with CuCl2 instead of the sulfate. So, take concentrated salt solution with a small amount of HCl to make it nicely clear. Immerse the Al-foil in this, and see that it does not react. Then put the solid copper salt on it and step back. Dissolving of Al in HCl or NaOH can be explained easily (oxide form Al(3+) and water with HCl and it forms aluminate AlO3(3-) and water with NaOH). But the fact that essentially non-acidic copper(II)/chloride solution is so effective remains very remarkable to me. With this Cu(II)/Cl(-) combo and Al-foil, you also have a nice and very cheap source of making lots of hydrogen gas! Fun assured!

I do not like this kind of questions. Do you want us to make your home work? I'll give you some hints, but please, next time, show that you've put some effort in the questions yourself. Compute volume of 1 mm of fiber. I assume you know how to compute the volume of a cylinder with known height and radius. Volume*density--> mass. MW(SiO2) gives mass per mole of molecules. Now you can compute total number of molecules of SiO2 and that gives you the number of O-atoms. FA? You mean force of attraction? Radius r is much larger than size of ions. This allows point-charge computations. Use F=k*c1*c2/(r*r), where c1 and c2 are charge and k is a well-known constant (I assume you know that or can find out). Quantum numbers are as follows: principal quantum number, which determines the size of the orbital (the shell number). angular quantum number, which determines how the orbital looks like (0 = spherical, s; 1 = two lobes, p; 2 = four lobes, d; 3 = even more complex , f). The larger the shell, the more possible shapes there are for orbitals. Please lookup yourself what relation there is (it is simple, I promise you ) magnetic quantum number, which tells you how an orbital is oriented in space. A spherical orbital can only have one magnetic quantum number, it does not matter how it is oriented and hence, each shell has just one s-orbital, but a two-lobes orbital can have three different orientations, perpendicular to each other, so there are three possible p-orbitals per shell. Finally, electrons have spin, which can be -1/2 or +1/2. In one orbital two electrons can exist, but only of opposite spin. Now it is your challenge to enumerate all 11 electrons in sodium, telling their principal quantum number, angular quantum number and magnetic quantum number. The spin of all electrons also can be determined, except for one. Beware, for one electron you must make the assumption that the atom is isolated, otherwise (in real metallic sodium), you have one electron per atom in a conduction band and cannot simply speak of an atomic orbital. It is up to you to decide which atom is doing such nasty things . All other 10 electrons are nice to you and are well-behaved . Use electrostatic formula, which I have given before. Now you can compute the distance between the two point charges. If you are not allowed to make the simplification that the ions are point charges, then things become MUCH more complex. I cannot deduce from here, however, whether you need to perform that more complex computation. Using the distance between the two point charges and the radius of the S(2-) ion, you can compute the radius of the Mg(2+) ion. A nice geometry question. Proceed as follows: How many atoms are there in one lattice cell? You know the size of a single lattice cell and you know it is cubic, so you can compute the volume of a single unit cell. Now you can compute how many unit cells there are in 1 cm3 and from that you know how many atoms there are in 1 cm3. You know there is 16.6 grams per cm3, so you can compute the mass of one atom and you can convert that to an atomic mass number. This is a matter of plain geometry. You know the size of the atom, the c/a ratio tells you something about the size of the cell (related to the atom's size). Lookup, how the atoms are arranged in a HCP crystal structure and determine, what the unit cell looks like. Computing the volume of that object is plain mathematics and hardly has anything to do with chemistry. From the density you can compute how many atoms there are per cm3. Take the third root and you have the average number of atoms along 1 cm. Now imagine as if they are packed FCC or BCC. Does it fit or do you need overlapping spheres for the atoms. One structure does fit, the other doesn't.

Best is to use a pure iron nail. If there is some zinc in it, that does not bother. After electrolysis you can easily get rid of the zinc hydroxide by dissolving it in excess sodium hydroxide. Most other metal impurities also result in formation of impure iron hydroxide. Why not go to an electronics shop, where they sell electronic components? At such places they also sell almost pure FeCl3.6H2O for etching purposes. From that you can easily make a lot of fairly pure Fe(OH)3 by simply adding solution of sodium hydroxide to a solution of FeCl3 in water. The product Fe(OH)3 will contain a little chloride, but that does no real harm. On heating it will be driven off as HCl.

Chemistry is so much more than watching things blow up... What are you interested in, in chemistry or in things going KABOOM?

@xeluc: Fe(OH)2 does not decompose, it is very sensitive to earial oxidation. A little bit like the problems you have with making nice CuCl. Have a look at this: http://woelen.scheikunde.net/science/chem/solutions/fe.html Look at the section for oxidation state +2 and quite some info is given on Fe(OH)2. Also look at the nice light green color of pure Fe(OH)2 and the color of the partially oxidized stuff. @airkyd: The materials you use are impure. You have an iron nail and copper wire immersed in the liquid at the same time. As Xeluc already stated, you must assure that ONLY iron is touching the liquid. So, wrap some copper wire around the head of the nail and only dip the lower part of the nail into the liquid. Then start over again and show us again what kind of precipitates you obtain.

Molecular weight: Roughly speaking: MW(CH4) = 12 + 4*1 = 16 MW(C2H6) = 2*12 + 6*1 = 30

Indeed, my most potent oxidizer is potassium persulfate. It is even capable of oxidizing silver ions to the +3 oxidation state! Its redox potential for the half reaction S2O8(2-) + 2e ---> 2SO4(2-) is over 2 Volts, which is REALLY high. On the other hand, this oxidizer acts very sluggishly. It is powerful, yet slow.

The reason that I used acetone and ether rinses was to make the stuff dry as fast as possible. Once it is thoroughly dry, it is not so sensitive to aerial oxidation anymore. But even the tiniest traces of water make it very vulnerable and that is why I did not succeed in making a sufficiently pure sample. Even under acetone it changes color . CuCl is somewhat soluble and the dissolved matter is what is oxidized so easily. I think it has to do with the larger contact area of the dissolved matter and the fact that the dissolved matter is complexed: CuCl2(-) is so sensitive towards aerial oxidation.... Your idea of boiling away water from a sample in a loosely capped container may work. I personally never tried that. I'll try such a thing at a test tube scale and I'll let you know. I can imagine, however, that at higher temperatures the CuCl hydrolyses and that it becomes contaminated with oxide and hydroxide ions. My aim indeed is a purely dry sample. I do not like wet samples of chemicals. Look at my sample of vanadyl sulfate on my website, it has a really nice color, but still it is ugly with all the water around. One of my chemistry books states the following way of preparation: If I read this, then I think that for the home chemist this hardly is feasible. Making all your solvents and acids free of oxygen is not easy. You can boil them for a while to drive off oxygen, but on cooling down, they readily absorb some oxygen from the air again.

There are two explanations: 1) A lower hydroxide of iron is formed. Iron (II) hydroxide is green/white, iron (III) hydroxide is orange/brown. 2) The nail does not consist of pure iron, but another metal is involved. When the iron is used up, then the other metal goes in solution. Many metal hydroxides are white. Due to impurities this may look off-white.

We have lots of pure aluminium around us. The cans and foil you know are not the oxide but the (almost) pure metal.

CuCl does not react with water. If you keep out oxygen rigorously, then it can be kept indefinitely, either dry or moist. The compound itself is stable, the pain is just in its extreme sensitivity towards oxygen, when it is wet or moist. I recognize very much of your struggling with the stuff . I've done this many times myself and I could make beautifully white stuff, but as soon as I tried to make it dry, it became greenish/brown and it was spoiled within seconds. Finally I've given up and I bought some of the stuff (now approximately 1 year ago). The purchased stuff keeps fairly well and it does not degrade further. The stuff is not hygroscopic like CuCl2. I even have tried to use acetone for cleaning up. I decanted most of the water, added acetone to it (which dissolves most of the water and remains of acid), decanted the acetone again, added new fresh acetone and decanted that and did a final rinse with diethyl-ether. But even that method of drying did not give me a nice white powder. If you really want nice white CuCl, then you have to work in an all-inert gas apparatus. That, however, is beyond the possibilities of what a home chemist can afford. So, I've given up on this . @RyanJ: I purchased some, as you have seen on the picture. It remains like this and does not deteriorate further. So, you can buy some of this stuff, but do not expect it to be 100% pure when you receive it. It will contain quite some contamination with copper (II). If you want to perform experiments with CuCl, then you can also make it in a wet environment. Xeluc has done so and indeed you can obtain snow-white stuff as long as it is kept under water. I think it is best that Xeluc gives details on how he made his CuCl.

I do not believe this works. Enantiomers have exactly the same physical and chemical properties, except when they are brough in contact with other enantiomers and when the change of direction of polarization is involved. In living animals, plants and humans, of many compounds only one enantiomer is used, but chemically speaking enantiomers cannot be distinguished. The solubility in any non-optically active solvent will also be the same for two enantiomers.

Knowing only the mass is not sufficient. If you also know the volume (at a known pressure), then you can determine the ratio by solving two equations with two unknowns. If you know the mass, then you have one equation: x MW(CH4) + y MW(C2H6) = m If you know the volume at a certain pressure, then you know the total number of moles of molecules (e.g. at standard pressure and 298K one mole of an ideal gas takes approximately 22.4 liters of volume). Let's assume there are A moles of gas. Then you have a second equation: x + y = A With basic algebra, you can solve this set of equations, giving you x and y as number of moles of each compound.

In fact, xeluc, you can't without really good lab facilities. Storing under oil is a real pain. Sodium comes in nice large lumps, CuCl is a powder. If you dry it, it will become darker already. Adding oil to the wet stuff will be a REALLY messy crap. Have a look at my site. I purchased some CuCl from a good chemical company and this is the best they could offer me. http://woelen.scheikunde.net/science/chem/compounds/cuprous_chloride.html If chemical companies are selling this as the best they have, then forget about making a purer compound at home. Even this sample was a rip off already. Making CuCl of this purity already costs almost $1 per gram on a commercial scale.

NO2 readily dimerizes, but the dimer also readily decomposes back to NO2. This is a large difference with e.g. CH3. Dimerized CH3 (ethane) never will form CH3 again, with NO2 there is an equilibrium: 2NO2 <---> N2O4 Low temperature and high pressure drive this reaction to the right. This is nicely demonstrated if you have a syringe, filled with NO2. If you press a capped syringe with the gas, such that pressure increases, then you'll see the color become lighter (after an initial transient, where it is darker). If you then suddenly release the thing, then you'll see darkening instead of lightening of the color. Very nice experiment, I have done that once. I made my NO2 by first making NO-gas from acidified ferrous sulfate and sodium nitrite, pressing the gas into another syringe through a thin tube (water spoils the demo). I next made O2 from H2O2 and KI and pressed the O2 gas also into the syringe with NO, approximately the same volume. Then you are ready to play. Be careful though, NO and NO2 are insidiously toxic. Do this outside!

CO slowly oxidizes to CO2 in the atmosphere, but this is quite different from NO. NO is oxidized at once. If you have a syringe of colorless NO and you press the gas out of the syringe, then you'll see a brown plume of NO2 and all of it is oxidized within seconds. CO takes days or even weeks to be oxidized and that reaction also requires sunlight. CO is not a radical. It only has paired electrons.

Yes, there is something "weird" going on. Usually, all atoms in a molecule want to have filled shells. If you look at water, then the oxygen shares its two electrons with the hydrogens, but the electrons of the hydrogens also are shared. This makes all atoms in water happy. In some compounds, some atoms are not happy at all. Look at a compound like methyl, CH3. This molecule contains a very unhappy C-atom, with just 7 electrons. It would be much better if a fourth H-atom comes into play, but two CH3-molecules also can become happy if they join together, forming H3C-CH3. Such unhappy atoms are missing some electrons. In more scientific terms, such atoms have unpaired electrons. Molecules, containing such atoms are called radicals. So, CH3 is a radical and it is VERY reactive. Now the surprising point is coming. NO and NO2 also are radicals. The nitrogen atom in both molecules contain an unpaired atom. NO and NO2 are very reactive molecules, but in terms of radicals they are surprisingly inert, such that it is even possible to store them. This is impossible with CH3. That would react at once with itself, forming ethane (H3C-CH3) and with the container, in which it is stored. Another example of an 'inert' radical is ClO2. This bright yellow gas can be created in macroscopic quantities and can be stored. It still, however, is very reactive and easily explodes, but in terms of radicals it is remarkably inert.

There is no theoretical limit on the length of C-C chains. Real-life compounds, such as certain plastics and resins can have chains with lengths of many millions or even billions. The limits are of a practical nature. Every compound, made by humans, even the purest ones, has defects, such as impurities, irregularities in the crystal lattice. This is not different for long polymeric chains and hence in practice, the length is limited.

You won't get barium peroxide but barium oxide. BaCO3 --> BaO + CO2