UC

If you place your planet orbiting a red dwarf, you need to take into account that most of the radiation from the star is going to be infrared, which is relatively low energy and perhaps unsuitable for photosynthesis. Since most visible light will be red, a plant using red light will appear black. Purple pigmentation would indicate that there is purple and blue light around to reflect, which isn't the case with a red dwarf.

I take it that you are learning VESPR theory? There are two "shapes" in a methanol molecule, since the hydrogens don't have any p-orbitals to hybridize. One of the two is tetrahedral. Keep in mind that it isn't a single bond that is shaped, but the cluser of bonds attached to a central atom. You should look into lone pairs and how they situate themselves in hybridized orbitals There is a key difference between boron and nitrogen and it has everything to do with valence electrons. good luck

What kind of purity? I would be interested if it's semiconductor-ish grade (5 or 6 nines) Is it fine powder or small chips that still show the luster of large pieces? Just like silicon, it has rather dull chemistry unless you want to experiment with organogermanium compounds (for which the tetrachloride is a much better starting point). I would agree with YT on this one.

I fail to believe you did anything at all to search for this other than immediately come here and post. Allow me to demontrate: http://letmegooglethatforyou.com/?q=gundanium Try the entire first page of results.

The feeling I'm getting is that you started this thread with the sole purpose of rambling about that link in your sig. that was already there when you started the thread. I don't particularly appreciate this.

It's from an anime show, Gundam Wing....

Preparation of Chlorates by Disproportionation of Calcium Hypochlorite in Aqueous Solution Aside from their famous (perhaps infamous is more appropriate) use in pyrotechnics, chlorates have a fairly rich chemistry and make a valuable addition to the chemical inventory of the amateur experimenter. Iodic acid and iodates can be prepared by the reaction of iodine with solutions of chlorate acidified by nitric acid. Reduction of chlorates with hydrochloric acid generates ClO2 (generally mixed with varying amounts of Cl2 depending on stoichiometry), which makes for an interesting small scale demonstration (Do not scale up!). They can be used in place of nitrates as the oxidizer for the synthesis of chromates and permanganates by molten salt fusions and they can serve as heat boosting agents in exotic metal thermites. Perhaps most useful to the amateur chemist, however, is that chlorates can be catalytically decomposed with heating and small quantities of MnO2 to provide an ever ready and cheap source of oxygen gas without the need for a pressurized cylinder. The industrial preparation and essentially the only viable option for large scale preparation of chlorates is by electrolysis of sodium or potassium chloride solutions. For microscale and qualitative experiments, however, building a chlorate cell is time consuming and downright overkill for the quantities desired. An alternative procedure for the amateur chemist involves boiling down bleach, a rather odorous and time consuming affair, which tends to leave glassware etched for fairly small quantities of chlorate. Here, I present a simple and clean procedure for preparing potassium chlorate in small quantities from cheap and widely available calcium hypochlorite pool shock. Safety It should be more than apparent to anyone with even limited knowledge of chlorates that these compounds are dangerous. As a strong oxidizing agent, contamination of chlorates by organic material or other reducing agents can lead to fires or explosions. Spills of solutions soaked into wood or paper can later on accelerate fires. Acidification of chlorates, especially by hydrochloric acid (or by co-acidification with chlorides), can lead to the evolution of chlorine and chlorine dioxide gas. The former is a powerful respiratory irritant and inhalation may cause pulmonary edema. The latter is capable of spontaneous detonation when it reaches certain concentrations in air and is a severe respiratory irritant in lower concentrations. Contact of molten chlorates with organic material can lead to explosions. Chlorates are highly toxic by ingestion and are skin, eye, and respiratory tract irritants. They are used as nonspecific defoliants and weed killers, and are thus highly toxic to plant life. I urge you, at the bare minimum, to use gloves when working with any chlorates, and preferably goggles and a lab coat as well. As for the procedure in general, you are working with a close analog of bleach, and spills can and will damage or destroy clothing in the same way that bleach spills will. Contact with any part of your body should be avoided and the entire procedure should be carried out in a location with good ventilation as small amounts of chlorine gas are inevitably evolved during the procedure. Legalese Manufacture of any pyrotechnic device or explosive with the product of this procedure may or may not be legal, depending on your location and licensing. The manufacture of a pyrotechnics oxidizer (the completion of this procedure) also may or may not be legal in your area. All preparations are carried out at your own risk with respect to personal safety and legality. Neither Science Forums, nor the author will be held responsible for any and all damages that may be incurred as a result of following this procedure. By undertaking this experiment, you understand the risks associated with it (including any which may not be explicitly stated in the safety section) and agree to the above terms. Required Chemicals: Calcium hypochlorite (usually available with around 50% active material as pool shock) Distilled water Potassium chloride An acid (anything even moderately acidic will work) A reducing agent: a sulfite, bisulfite, metabisulfite (pyrosulfite), dithionite (hydrosulfite), or thiosulfate (hyposulfite) are all acceptable, with the first three being preferred. Equipment: Heat source sufficient to boil water Stir rod Beaker(s) and watch glass Filtering Setup (vacuum filtration greatly preferred) Ice/brine bath Experimental Into a 1 liter beaker containing 500mL of distilled water, 100.0g of 54.6% calcium hypochlorite pool shock was slowly added with stirring. The hypochlorite foams somewhat as it is added (hence the stirring) and small amounts of chlorine or chlorine oxides are evolved, giving rise to the familiar pool chemical smell. The liquid is a pale green-yellow, the same color as bleach, indicating dissolved hypochlorite ion. The beaker was covered with the watch glass to exclude as much air as possible. Heating was carried out in a microwave oven on maximum power. The mixture was watched at all times as it is highly prone to “boiling” over, especially at the beginning. The solution was brought to a hard boil and was kept boiling for 20 minutes. During this time the solution reduced to about half of its initial volume and the small pellets of hypochlorite broke down into a pasty suspension consisting (most likely) of calcium hydroxide. The smell of chlorine is detectable for the majority of the heating time. The solution was vacuum filtered after cooling. Around 235mL of clear filtrate was collected. The solids were relatively dry and were discarded. The filtrate was transferred to a 600mL beaker and 19.0g of potassium chloride was added. This was covered with the watch glass and boiled until the solution volume was about 130mL. During the boiling, the solution became cloudy, likely due to the reaction of some traces of calcium hypochlorite with carbon dioxide in the air. The solution was gravity filtered while hot through a loose cotton plug to remove the cloudiness and placed in a salted ice bath. Potassium chlorate separated as flat, colorless, glittering plates which appeared more feathery in form when cooling was rapid. The crystals were brought into suspension by mixing with a stir rod and were vacuum filtered. They were washed with two 25mL portions of ice cold distilled water and air was drawn over them for several minutes to dry them as much as possible. They were dried by leaving at room temperature for about 24 hours (the solid is not at all hygroscopic and readily dries without the need for a dessicator) to afford 15.2g of odorless, free-flowing, glittering plates. This was 48.7% of the theoretical yield based on calcium hypochlorite. Waste Disposal As stated in the safety section, chlorates are very toxic to all forms of life. They are also a very persistant drinking water contaminant. To prepare the liquid wastes for disposal, a spatula of the chosen reducing agent is added and acid is dripped in. The products of acidifying chlorates; chloric acid, chlorine dioxide, and chlorine, and readily reduced by free sulfur dioxide released in the acid environment. The waste is very carefully smelled. If sulfur dioxide is detectable, all chlorate in the waste has been reduced to chloride. If not, more reducing agent and acid are added. Afterwards, the waste can be flushed down the drain with no toxic effects. Discussion The reaction of core importance in this experiment is a disproportionation reaction. In this kind of reaction, one reactant is both oxidized and reduced. For this experiment, the reactant in question is the hypochlorite ion, which reacts as follows: 3ClO- => 2Cl- + ClO3- The hypochlorite ion is stabilized in the presence of base and by low temperatures. Boiling the solution provides the necessary energy to drive the reaction to the right. The competing reaction is the decomposition of calcium hypochlorite: Ca(OCl)2 => CaCl2 + O2 This reaction is notably observed when heating dry calcium hypochlorite. No chlorate is produced, only oxygen gas. The calcium hypochlorite in pool shock is probably prepared by exposing damp calcium hydroxide to chlorine gas, then mixing in some extra hydroxide to stabilize it. The product of this preparation is probably a mixture of calcium hypochlorite and basic calcium hypochlorite (Ca(OH)(OCl)), the latter which is probably only weakly soluble in water. On the scale of a swimming pool, it will dissolve, but not in 500mL of water. It is likely that this undissolved solid undergoes the above decomposition reaction when boiled instead of the disproportionation reaction. The evolution of oxygen is probably why the solution seems to “boil” when heating starts and why it seems very vigorous. The label of the pool shock, however need not distinguish between the two compounds as both will work in a pool. The meager 48.7% yield is thus likely a misrepresentation, as it should be based only on Ca(OCl)2 content. Calcium chloride and chlorate are both extremely soluble. Potassium chloride is fairly soluble, but potassium chlorate is only weakly soluble at low temperatures. When the solution containing potassium ions, calcium ions, chloride ions, and chlorate ions is cooled, the least soluble combination is potassium chlorate, which crystallizes out. This is an example of a metathesis reaction (double displacement). 2KCl (aq) + Ca(ClO3)2 (aq) => CaCl2 (aq) + KClO3 (s) Another reaction relevant to the procedure is the reaction of calcium hypochlorite with carbon dioxide and then the reaction with chlorides in solution: Ca(OCl)2 + CO2 + H2O => CaCO3 + 2HOCl HOCl + Cl- => Cl2 + OH- This produces free hypochlorous acid, which is highly unstable. This may then react with chlorides, allowing chlorine to escape solution. This is especially relevant for calcium chloride, since the solution is already saturated with calcium hydroxide. The precipitation of Ca(OH)2 drives the release of chlorine, which may escape (especially when it forms near the surface, where CO2 is likely to be, and especially because the solution is boiling) before reacting with the excess of Ca(OH)2 in solution. The beakers used to heat the solutions in this experiment end up with a thin film of calcium carbonate on them from this reaction. It appears to be etching at first, but can be removed by vigorous scrubbing or by acids. Aside from stopping splatters, the watch glass is an attempt to minimize how much CO2 gets into the solution. References and various links: Initial discussion on sciencemadness.org: http://www.sciencemadness.org/talk/viewthread.php?fid=2&tid=10727&action=printable Preparation of chromates with chlorates: http://webpages.charter.net/dawill/tmoranwms/Chem_Chromate.html MSDS for potassium chlorate: http://www.jtbaker.com/msds/englishhtml/p5620.htm What not to do with chlorates: http://www.destructve.com/bromicacid/mistakes.htm#4 Chlorine dioxide demonstration by Woelen: http://81.207.88.128/science/chem/exps/clo2/index.html Destruction of Chlorates: http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/destroy.html

Even if it works you'll find that calcium sulfate is surprisingly and annoyingly soluble in nitric acid, increasing in solubility with HNO3 concentration.

Sounds good Phi some exotic flavors for you: spicy curry fillet of panda with wasabi road apple with pecans and cinnamon ethidium bromide (follow your DNA while you eat!) strawberry cesium butanethiol berry triethylamine swirl A nice list can also be found here. Just make the candies into ice cream flavors:

Keep in mind that Li2O is an extremely hygroscopic, stripping water from the air to make Lithium hydroxide. Lithium oxide also has a very, very high melting point, above the boiling point of lithium. I suspect that he wouldn't have collected much metal if it was being evolved as a gas that would ignite on exposure to air. Also, it may well be that lithium oxide is soluble in molten lithium hydroxide and what they thought was the oxide was a mixture of the two. Lithium hydroxide has a sufficiently low melting point for Lithium to be formed as a liquid.

One has to consider the multiple definitions of acid in this case, but since we are discussing a hydrogen halide anyway, any will be appropriate. The fluorine, being highly electronegative in comparison to the hydrogen, pries the electrons away from the hydrogen. This is now an electron deficient hydrogen, which satisfies the lewis acid definition (which is the most useful definition once you get into organic chemistry). For the more well known arrhenius acid definition, this electron deficient hydrogen is fairly happy to break its strained bond to fluorine and attach itself the the lone pairs of electrons on the oxygen of water, forming a hydronium ion.

This should be in homework help, but use this: http://www.cem.msu.edu/~reusch/VirtualText/aldket1.htm and scroll down to acetal formation. That online text is very useful for many things. Also, download this program. Your drawings will improve immensely http://www.acdlabs.com/download/chemsk.html For the intermediate step in the middle row, you should draw the carbocationic tautomer, which is the reactive species. Only use one arrow, and it should be obvious that the oxygen in the hydroxy group doing the attacking. For the bottom row, water will not be doing the deprotonating since acetals are made in conditions that remove water, but with acid catalyst. You will see why water removal is necessary when you complete the mechanism. The thing that deprotonates it will be the anion of the H+ earlier in the mechanism. In this way, the acid is regenerated and is a catalyst, not a reactant. I have drawn up the mechanism and will post when I see your full attempt at the mechanism.

The standard electrode potentials are for, I believe, 1M solution of the metal compound as a neutral solution in water. The values are not set in stone and change based on pH, concentration, solvent, etc. Plating baths are almost never just a metal salt in water. They have all sorts of additives to get the potential below that of hydrogen and to ensure that the metal adheres properly and with an appropriate surface finish.

If you follow the link in the lithium article to the Sir Humphrey Davy article, you will find that Lithium oxide is almost surely a typo and should be lithium hydroxide. Electrolysis of molten lithium hydroxide generates the molten metal at the cathode

Yes, all of those compounds are salts. A salt, in the chemistry sense, is an ionic compound that can be formed by the neutralization of an acid by a base. Some of these are not soluble in water, but those that are dissacociate into ions in solution. They are also known as electrolytes. The common use of salt is sodium chloride which fits this description. Sodium tetraborate (borax) and various sodium acid salts of ethylenediaminetetraacetic acid are also used for laundry, but the chemistry sense of "salts" covers an essentially infinite number of compounds.

I realize this. It does get rather sketchy in places, but you can't argue atomic number or the s, p, d, and f blocks as a general means of organization. The properties of the elements arise at least partly, if not mostly from their electron configurations, and as a result, are just as inconsistent in places as the electron configurations.

Well, if you want to get technical, it follows the number of protons in the nuclei (and hence order of filling of the ellectrons) and the quantum states that the electrons are in: aka: shells and subshells.

No. The 12 volt Lithium battery you refer to is probably actually a Li ion battery. These batteries IIRC, rely on cobalt redox to generate current and use a lithium salt as a carrier. Smaller batteries that are not meant to be rechargable are that way for a reason. Reversing current through them is either highly inefficient at regenerating the compounds that give you the current, or may lead to gas evolution and rupturing the battery, perhaps with some violence. AA type lithium batteries actually use lithium metal and are not rechargable.

Sedit- check out the Cannizzaro reaction. Base around benzaldehyde is a no-no

What size coil are you looking to make and out of what material? As alien said, a drill and some rod stock (metal is much better than wood. in some cases wood will be useless) is your friend for narrow coils. For larger ones, I recommend a bit of pipe locked into a vice and some hand power. I make coils all the time, then cut them up one side to make rings for chainmaille armor and jewelry.

Spiffy Merged post follows: Consecutive posts mergedand at the same time, delicious. Very nice! It's a solution for those who are incapable of using the edit button. Speaking of which, is therea time limit for editing one's posts? It appears that there is, but how long is it before they lock? Merged post follows: Consecutive posts mergedAlso, I love waffles

Ibeamer, you might also want to note that the reaction gets hot, extremely hot if you use strong lye and lots of aluminum foil. That is, however, a decent way to get hydrogen. DIY Electrophilic aromatic substitution: Required: -Test tube or similar small, clear container -Dropper -Sodium bromide (NaBr. This is available as a bromine reserve starter for hot tubs/ spas) -An acid (HCl, H2SO4, H3PO4, or any reasonable strength acid is acceptable. Sodium bisulfate is marginally acceptable, but stronger acids produce better results) -Hydrogen peroxide solution (3% is available in most food stores as an antiseptic. Other concentrations are available if you know where to look) -Chloraseptic spray (I'm not sure if this is available outside of the US. Make sure the bottle says it has phenol in it) A tiny spatula of NaBr is placed in a test tube or any clear container that you can observe the reaction in. Add a few drops of hydrogen peroxide solution. I used 6.5%, but 3% will work as well. Just use a little bit more if you're using 3%. The solution is acidified. I used a few drops of concentrated HCl, which produced a red-brown solution of bromine in water. Even sodium bisulfate (with a little water) will work, but the mixture appears yellow because the concentration of bromine is lower. The mixture is diluted a bit with water, so there is some room to see the reaction above any undissolved solids, and a few drops of chloraseptic spray is added. The chloraseptic spray is a source of phenol, a good example of an activated aromatic ring. The color of the solution fades as the bromine reacts, and a white precipitate forms. This precipitate is 2,4,6-tribromophenol, which has an intense antiseptic smell. With the red-brown solution, the drops of chloraseptic formed a thick white layer upon hitting the bromine water. With the dilute, yellow solution, there was a bit of delay before much precipitate formed, which clouded the solution, instead of forming a compact glob. The quantities used here are extremely small, and the few milliliters of waste (containing at most a few milligrams of tribromophenol) can be disposed of in the trash. Happy experimenting

I yell at the screen sometimes when I see stupidity...like the mythbusters episode where they put an entire pair of blue jeans in a huge bomex (!!! can't they afford pyrex or kimax) beaker full of H2SO4 and HNO3 (They appeared to be reagent grade which further confuses me as to why they were using cheap bomex stuff, which is thin walled and often has bubbles in it). I haven't seen the episode in a while, but I don't think there was an ice bath, so the reaction ran away and started spewing NO2 fumes. Also on the hindenberg-thermite episode, they failed to test all combinations of experimental conditions. It was really glaring to me, but probably not to most people. At any rate, the show is a lot better than brainiac, despite a lot of little mistakes and a few big ones.

I'd reserve a retort for high boiling materials that demand strong heating (and will condense easily with only air cooling), but that aside, cooling the neck in some way will give you your desired effect. You could run ice water over it, but you'd need to attach some sort of something to the end of the neck to keep the water from getting in your distillate. A bead of caulk intended for bathroom work would be a good waterproof "glue" to attach such a barrier