UC

A list of highly endothermic dissolutions (a little more than halfway down): http://chemieunterricht.de/dc2/tip/08_98.htm From my understanding, Eis is ice (preferably as snow) and kaltwasser is "cold water" (ice cold, I can only imagine). These temperatures are probably only available if you use finely powdered solids and snow and mix the two rapidly. I have 5kg of MgCl2*6H2O that I bought on clearance for $1 at walmart last year. Sadly, it's a little late for snow, but I can try with finely pulverized ice.

I have a copy of Zumdahl's chemical principles 5th edition. I take it that you are using a similar book. Page 329-330 in the 5th edition. (I hope you know where it is in your text. Otherwise, use the index) Qualitative analysis. Are you using periodic table "groups" or the "groups" given in the book? I-insoluble chlorides, II-sulfides insoluble in acid solution, III-sulfides insoluble in basic solution, IV-insoluble carbonates, V-alkali metal and ammonium ions. The two failed tests in question #1 exclude which two groups? The precipitate formed in the problem identifies which group? Are the remaining groups even able to form this kind of precipitate? Your answer to #2 looks right. There are ways to seperate the two, but those are outside of what you need to know for your course, I believe. Adding perchloric acid, and chilling the solution will precipitate almost all the K+, for examaple. The difference in solubility is quite large. Oxalic acid will do the same, but with less selectivity. The potassium salt is about 10 times more soluble than the sodium salt. If you used one and then the other, you could achieve clean seperation with only a small amount of the original salts lost.

Not with very selective absorption filters neon doesn't Take a peek at the spectra sometime (ask the physics dept. They probably have the right piece of equipment. the name escapes me) Granted, AFAIK, they are all red, yellow, and orange, but still Very cool stuff.

http://www.cem.msu.edu/~reusch/VirtualText/benzrx1.htm Everything you need is on that page. Feel free to ask if you need it explained better.

That's the single most annoying thing I've heard in a while. A forum is generally where to ask stuff that you can't figure out elsewhere or need better insight that the internet fails to provide. You're wasting everyone's time. Also, you could have typed in everything you need to get to that wikipedia article in less words than it took to make that post.

1: Sulfuric Acid- undisputed king of acids in importance. Now if only I could get some that isn't brownish and sold as drain cleaner.... 2: Mercury- *shakes it around it's little bottle. squee!!!* I have a very unhealthy love of toxic things and mercury is way up there. Plus it's so cool looking, makes good thermometers work, and is used for ever-useful amalgams. 3: TCCA- trichloroisocyanuric acid. The possibilities are endless with this close relative of N-chlorosuccinimide. Not the least of which are radical chlorination, chlorine generation, haloform reaction en-mass, oxidation of amines to nitriles, oxidative decarboxylation of alpha-amino acids, ether oxidation (to esters), etc. Plus you can buy it by the bucketfull in lots of stores. 4: Thiourea Dioxide- A very powerful over-the-counter available reducing agent. Not much has been done with it yet, but I believe it will be recognized as a must-have for amateur chemistry in due time due to it’s (claimed) versatility. In the meantime, some pretty blue shades can be had by using it to reduce indigo dye powder. The ability to go from ketones to alcohols (with no hydrides anywhere around) in high yields opens a ton of doors (and holds promise for lots of other transformations) 5: Chloroform- Something I made early on in my experimentation. Smells really good and sweet. So far, other chlorinated hyrocarbons are all disappointing smelling. Good solvent and who doesn’t love CDCl3 for NMR? 6: Nitric Acid- I’ve got none, but it’s got a list of uses a mile long. One I’ve been meaning to try is paracetamol  p-benzoquinone (or p-chloranil depending on how much HNO3 you use) via aqua regia oxidation. Can dissolve just about anything when paired with HCl. Aromatic nitro compounds are also integral building blocks for larger molecules, being precursors to aniline/azo/halide/cyano/phenol groups via reduction/diazotization/sandmeyer reactions. Plus picric acid is so incredibly yellow and I want a swatch of cloth dyed with it. 7: 6,6-dibromoindigo- easily the most expensive dye in the history of human civilization; tyrian purple. This is the one Alexander the great had a whole outfit dyed in just to show off how rich and powerful he was. One of the first chemicals I ever wanted to synthesize and I have plans for it lined up this summer. 8: Iodine- smells great except up close. Beautiful as gas. Excellent as a leaving group. Very pretty compounds like lead iodide. Lots of synthetic utility. 9: Platinum- ‘nuff said. 10: (E)-(4-hydroxy-3-methoxybenzyl)-8-methylnon-6-enamide – mmmm. Spicy.

Funny, I have a pound of it in the drawer to my left from the very company...and 2lbs Zn, 1lb Cd, 4lbs Pb, 1lb Sb, 1lb Sn, plus some lead and zinc wire. I see they've wisely stopped sawing the cadmium sticks into pieces. (I have half a stick) The bismuth is only $15/lb now, plus whatever shipping is. $1US seems to be about $2NZD right now for reference.

*puts money on a ban in the near future* drive by trolling much?

is there going to be a difference in the chemical shifts for the more substituted alkene versus the almost-symmetrical one? This database has everything you could possibly ever need as far as numbers. Both compounds and the chemical shifts are given under cyclohexene: http://www.chem.wisc.edu/areas/reich/Handouts/nmr-c13/cdata.htm Of course, the point of using C13 NMR is out the window if you can't understand why the shifts differ between the two (and I suspect that you will be required to explain).

Probably almost exactly the same.

BURN HER!!! (couldn't resist, sorry)

I was given that article/memo when I started undergraduate research as a safety warning. Not that we are working with tBuLi (although there is some in the lab freezer). The newer technologies use things like 5-amino-tetrazole. 5 nitrogens, 3 hydrogens, and a carbon (with a ring!!). Avoiding hot metal compound side-product is a good thing.

Sorry, I see that I wrote the formula for aluminum hydroxide wrong. I guess I was thinking of aluminum oxide at the time ([math] Al_2O_3[/math]). It should be [math] Al(OH)_3 [/math] Sulfuric acid is a good choice of electrolyte. Keep any halides, which readily corrode stainless steel, out of the reaction mix. Sea salt is identical to kosher salt or uniodized table salt. Sodium chloride is the only significant component.

bis-(2-ethylhexyl) phthalate (common PVC plasticizer) should work as well. Alkyl benzoates as well, I assume.

No it doesn't. You'll just evaporate a lot of iodine... Heating with concentrated nitric acid is the standard I believe. Bromates or chlorates dissolved in dilute H2SO4 or HNO3 also oxidize elemental iodine to iodates. Bubbling chlorine through I2 in water will give iodates and eventually periodates. Woelen has some experiments on this. It can also be done electrolytically, but I don't think this is as efficient as for bromates or chlorates.

Pale blue? That would indicate the presence of NO in your sample of NO2, forming N2O3 upon cooling. It is fairly dark blue when pure, so the concentration is probably quite low. I imagine that measurement error of some sort is to blame instead.

Do you mean thiocyanate complexation? That will only keep the iron dissolved. The thiocyanate is quite expensive too. John Cuthber has the right idea I think.

Oxalic acid simply doesn't work here, or a lot more people would have made this stuff by now. Quoting Alexein from his post on sciencemadness.org: "Start with 2,4,6-trichlorophenol and dissolve it in dry toluene or dry dichloromethane. "wet" solvents are ok but the yield will be reduced. Add one molar equivalent of a suitable organic base. I used triethylamine. stir until everything is dissolved. chill the solution with ice and add in SLOWLY 0.5 mol equivalents of oxalyl chloride. Its going to heat up.... ALOT... so keep it stirring and chilled. Once its all added remove the ice and let it warm to room temperature with continous stirring. I find purity increases if you let it stir overnight. The TCPO is a fine white powder mixed with triethylamine hydrohloride at this point. Suction filter the mixture and keep the suction on to dry out the powder. Discard the filtrate. pour water directly onto the white powder and wash out the triethylamine hydrochloride. Dry the powder again. The powder is TCPO, its still pretty impure but usable. If you want nice big pure crystals (like the ones in the picture) you can dissolve TCPO in tetrahydrofuran, filter, and recrystalize. "

The reaction is extremely exothermic (readily proceeds beyond the point where you can touch the reaction vessel) and produces a nice mist of lye solution along with the gas. Eventually, the steam being produced by the hot liquid thins the hydrogen out quite a bit. It does work well for the first few minutes though...and then you're left with caustic [math] Al_2(OH)_3 [/math] sludge to get rid of.

A capacitor is effectively a break in the circuit, AC is a constant shuffling back and forth of electrons in the wires. There is no push against the capacitor to build up any charge for the reasons buttacup stated above. You need to use DC current if you want anything to happen. Fundamentally, capacitors block DC current (and build charge) and pass AC current. Some kinds are unidirectional and blow up if you wire them backwards while others will work in either direction. You should use copper sheeting anyway. It's quite a bit sturdier

Species native to an area are generally in a very delicate balance. There are forests a short drive from my house that at one point had an understory entirely composed of magnolia. However, white tailed deer eat the seedlings and destroy saplings. I took a several mile walk through the woods and found a single stunted magnolia. Flocks of birds depended on the oily magnolia fruits for energy. Now, thorny, toxic barberry has taken over the open spaces. Nothing eats the plant itself, but the birds eat the fruit, which have virtually no nutritious value. The bird population has declined in the area and the barberry spread even more rapidly as a result of bird dispersion. Invasive species are invasive for a reason. They grow fast, spread rapidly, are often toxic or thorny, and have little food value to native species. Even tall maple trees in the area are not the natives. They are mostly norway maple now. These are resistant to diseases and produce huge quantities of viable seeds. They've wiped out the native hardwoods and grow thickly enough that the ground beneath them is literally bare except for a few japanese stiltgrass plants, which choke out anything growing on the ground elsewhere. These woods have not been cut since colonial times. There is an unfortunate clause in the deed (the land was donated to a local college) that prevents them from intervening to save the woods. The sentiment when it was written was that the woods will stay as is if left alone, but this was before many invasives were around, something the deed-writers could not have forseen. Now everything around the wooded area has been cut down and it's overrun with invasives. The woods are literally sick and it's depressing to see because nobody can do anything about it. We are losing ecological diversity and a piece of history. There's also something beautiful about a healthy forest. It's something you'd want around for future generations to be able to see. In this particular forest, it is also a valuable research tool. The invasives make it harder, but studying what's left of the natives helps archaeologists better understand the past. In this particular area, it is quite unusual that there is an old growth stand at all given the suburban development around it. These natives are the resources that prior generations had available to them in the area. Studying the invasives themselves can help, even though they can't be removed at this point. Better models of their growth and spread are available through study which will be valuable elsewhere. It's making the best of an unfortunate situation.

2. The iodine/starch titration is going to be the most accurate determination technique by a wide margin. Titration is a skill and if you haven't done it before, you should practice at least a few times.

London Dispersion Forces. Long hydrocarbon chains especially should give lots of space to interact between molecules. If you have tons of even weak bonds forming, the macroscopic results will be quite significant.

Reread the original post. Those are two seperate questions being asked. The stacking refers to the foil thickness only. How would that help to calculate "empty space" anyway? For this, examination of the lattice structure is necessary. The given mass and volume produce the theoretical density for aluminum, showing it is a "solid" metal cube with no macroscopic voids.

Essentially a silicon thermite then. It will take a lot of heat to start. Look at how close the electronegativites are as well. I think it would be more of an intermetallic/semimetallic than an actual silicide capable of forming silane upon acidification. You would still produce silicon, just mixed with alumina and possibly dissolved/alloyed with excess aluminum.