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Activated Carbon - Desorption in sunlight?
I have some activated charcoal pads or small "pillows" that can be used to adsorb odors and chemical fumes from different materials.
The instructions say to put them out in sunlight to reactivate them once they have been used for a while.
What is the mechanism whereby sunlight is able to reactivate the charcoal?
I tried googling but could not find a definitive answer. I ended up assuming it must be just heat that forces the gasses trapped in the charcoal to come loose and evaporate, thereby restoring some of the adsorption potential. Is that the mechanism or what is the sorcery going on here?

Your guess is right: the mechanism of charcoal regeneration is undoubtably temperature-induced desorption. Gas molecules are adsorbed on the large surface area of activated carbon by a mechanism of physisorption (or physical sorption) caused by weak nonspecific interactions (van der Waals forces) between gas molecules and the carbon surface. Unlike chemisorption, this process is reversible and, since it involves only weak interactions, it can actually be reversed by relatively small increases of temperature.
Wikipedia gives a specific example:

The activated carbon can adsorb a large amount of methanol vapours in ambient temperature and desorb it at a higher temperature (around 100 degrees Celsius). During the daytime, the sunshine irradiates the collector, so the collector is heated up and the methanol is desorbed from the activated carbon. In desorption, the liquid methanol adsorbed in the charcoal heats up and vaporizes.

You will surely notice the resemblance between this text and your case.

Would it be reasonable to assume that any heat source could regenerate charcoal?

@F'x The example from wikipedia seems to have disappeared. The exact paragraph is here. It also seems to show up verbatim in a bunch of online encyclopedias without a source, so I don't know where it was from originally.


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Choosing the right condenser
What are the advantages and disadvantages of the various types of condensers commonly found in the laboratory? Obviously more intricate pieces of glassware are more costly but assuming they are all available why would one use a Dimroth condenser vs. a Friedrichs when refluxing something? I've heard that Graham condensers are to be avoided when refluxing due to the possibility of clogging, and yet it is still very common - when is it appropriate to use one?

Illustration made with ChemDraw. Feel free to reference other designs but please attach a schematic for clarity.
Edit: This question was motivated by an organic preparation which involved the bromination of an alkene in boiling water. Bromine has a boiling point of 58.8 °C and on top of that the reaction was exothermic. It was difficult to avoid the loss of Br2 gas with your typical Allihn condenser, but the increased cooling capacity provided by a Friedrichs returned Br2 back to the flask as it was produced.
Also, while the Friedrichs forces vapors up a spiral path, the path itself is wide, and in my condenser at least, there was a bit of leeway for liquid to drop down the sides, helping to prevent blockages. I didn't try a Graham condenser but I imagine a much slower rate of addition would be supported by this condenser.

Also, I should mention that I've looked at the wiki page on condensers: http://en.wikipedia.org/wiki/Condenser_(laboratory), but I found it lacking in practical detail and there is little comparison between the different types.

So far answers to this question suggest that nobody really uses Graham condensers!  (Or that they aren't even condensers, just coolers.)

If you are considering cryogenic applications a Dewar condenser may also be appropriate.

@Mart 's comment impelled me to return to this question and correct my answer. I've deleted incorrect material and expanded the discussion to, hopefully, provide correct information. There is a good discussion (better than the reference previously cited) of the issue here.

Reflux is the process of boiling reactants while continually cooling the vapor returning it back to the flask as a liquid. It is used to heat a mixture for extended periods and at certain temperatures...A condenser is attached to the boiling flask, and cooling water is circulated to condense escaping vapors.

If you are refluxing a mixture, as you might in organic synthesis to increase the speed of the reaction by doing it at a higher temperature (i.e., the boiling point of the solvent), then any of the condensers that worked well enough to avoid the loss of solvent and avoid "flooding" would work equally well. When you're refluxing, you want the "reflux ring", the place where the vapor is visibly condensing into a liquid, to be no more than 1/3 of the way up the reflux column.
You have two different basic types of condensers shown, Graham-type condensers (the first 3) and coil condensers (the last two). In the coil condensers (the left condenser in the picture below), the water flows through the coil and the vapor moves up in the larger, outside  area of the condenser, condenses onto the cooled coils, then drips back into the pot. In a Graham-type condenser (the right condenser in the picture below), the water flows around a tube (whether straight or coiled) that contains the vapor/condensed liquid.(picture source) The Graham-type condensers clog (or flood) more easily since they have a more restricted path for the liquid to return to the pot.

Graham-type condensers: The Liebig condenser is simple, but has low cooling capacity and can be fairly easily clogged as the condensed liquid flows back into the flask and blocks the vapor that is trying to escape. The Allihn improves on this design by having a wider bore at the bottom and condensing the liquid on the "bubbles" where it can run down the sides and avoid blocking the vapor. (I've used this to good effect in refluxing many reactions.) The Graham condenser is the same basic design as the other two, but the condensation tube is coiled which provides more surface area for cooling...but also tends to send the condensed liquid right into the path of the vapor trying to move up. It is particularly prone to flooding.
Coil condensers, such as the Dimroth and Freidrichs, have high capacity for cooling with fewer problems from flooding since the vapor condenses on the coils and drips back from the little prominence at the bottom of the coils into the center of the pot. The vapor has an easy time getting past the drops falling into the pot. If you can afford it, this seems like a good choice for most applications. The Freidrichs condensers, which incorporate a cold-finger with the spiral, are higher capacity, quite bulky and heavy. I have seen them used with rotovaps where you are taking a lot of solvent off quickly, but not with an ordinary reflux apparatus. This would be over-kill for a simple reflux reaction situation.
Sorry for the incorrect information (for those of you who looked at this before) and hope this is helpful.

I'm not sure the "number of (theoretical) plates" concept applies to the condensor. In the condensor condensation energy is carried away by the coolant, a re-vaporation should hardly take place. The plates concept explains the rectification column. Important parameters for condensors are exchange area (for heat transfer), Volume (for retention time) and sucspetibility to clogging. However, I'm not a lab rat, my experience with that kind of equipment is very limited!

Good points...I messed this one up and have gone back and (hopefully) corrected my answer. Thanks for the heads-up. Next time I'll listen to that little voice that says "somethings not right in this".

@JaniceDelMar With flooding you mean that the vapor pressure of the boiling in situ reactants push the condensed vapor upward out of the condenser, leading to possible dangerous situations as well as reactant loss, because the reactants will flow out of your condenser?

While mostly right, this answer somewhat encourages the misuse of condensors and coolers. There are very clearly outlined use cases for each of the ones shown, and mixing them up is just bad laboratory practice, leading to failures and accidents.

Dimroth's is the only cooler really meant and generally suitable for reflux cooling.
The rest are meant to cool a product stream that only goes in one direction, or for destillation. Misused for reflux cooling, they all fail abruptly when too much condensate is produced, with the result of flooding.
Reason for this is the way downwards being too thin or even narrowing (Allihn, Friedrich, Graham).
The other problem is condensate flowing down on the walls, where it can dissolve your grease in the ground glass connections and sneak out, create concentration and temperature gradients in your reaction, start boiling already before reaching the vessel, possibly leading to foaming etc. With the Dimroth cooler, the condensate drops back safely directly into the vessel, onto the stirring bar, leading to a safe and homogeneous reaction mixture.
Even if the condensate contains some solid precipitate, it is washed back safely with a Dimroth cooler, while it will surely get stuck somewhere with all the others.
For the others:
Liebig diagonally downwards on top of a distillation column
Allihn more effective that Liebig, but only if it can be mounted vertically, obviously.
Friedrichs condense some very volatile distillation product
Graham only for a liquid (or noncondensing(!) gas) stream constantly flowing in one direction (heat exchanger).
simple coil condenser non condensing (or any downwards) gas stream

Soxhlet condenser
There is one more type of condenser, quite rare and often forgotten: the globe-shaped Soxhlet condenser. Typically used as a reflux condenser. The vapors pass between the outer wall, cooled by air, and the outer wall of the inner sphere filled with circulating coolant. Usually made of glass, sometimes metal (for better heat exchange, I presume).
The Soxhlet condenser is mainly used in the distillation of liquids with a high boiling point – the construction allows for effective cooling not just by one surface (as any "normal" condenser would work), but by two surfaces, including the air.
I've only seen it once in person, all the mentions on Google Books and Google Scholar are mainly dated somewhere between 1900s – 1920s, and as for the illustration, the only source I found is the Soviet textbook on lab techniques and glassware (arrows denote the  circulation of the cold water) [1, p. 61]

References

Voskresenskii P. L., Tekhnika laboratonykh rabot (Laboratory Technique); Khimia: Moscow, 1969. (in Russian)

Is this related to a Soxhlet extraction?

@Blaise Yes, in a sense that Soxhlet extractor can be equipped with the Soxhlet condenser. However, usually Allihn's or Liebig's condenser is used. Unfortunately, I'm not aware of whether historically Soxhlet extractor was designed to be used with the Soxhlet condenser.

I can't find interesting details on the practical use of the Graham column: I mean it may be more adapted to a reflux set-up or a distillation one but it looks difficult to get information regarding this.  Intuitively, I'm prone to think that the Graham design does not fit with the requirements of reflux heating: As mentioned earlier, it can be easily clogged by low-boiling materials, inner diameter of the spiral is so small that returned-dripping and upward vapors will circulate only with great difficulties.  I guess the best use for this condenser is for downward condensation in a distillation set up. Better again, avoid this old-fashion glassware: I can't remind the presence of a single Graham condenser in a chemistry lab!
For reflux boiling, there are 2 very efficient condensers: the Allinh is a good one for general applications, but the Dimroth & the double surface coil condenser can deal with a large amount of vapors and retain every single drop of solvent (if the appropriate length is selected wrt the volume).
The Friedrich is very useful in addition to a standard Liebig condenser, when placed at the end of the distillation setup its cooling surface makes it unavoidable for ethers, etc.
BTW: The Dimroth is the condenser of choice for soxlhet extractions, since the drip tip on the terminal loop of the spiral allows a constant dripping on longitudinal axis and straight to the extraction thimble.

Re-quoted misinformation leads to incorrect speculation. Here's what I've learned through practical experience over the years.

Not recommending it but if need be, Dimroth's & Allihn condensers (as well as other types) can also be stacked.
To reduce flooding, use larger joint sizes.
I assure you: a 45/50, 150 cm, heavy walled Allihn works fabulously for refluxing 11 Liters @ 120 C for 14 Hours.
Contra-wise, given the same parameters, you can't use a 24/40 orifice to reflux even 2.5 liters. Be it with a Dimroth, Friedrichs, or even a Cold Finger with Dry Ice. Whatever I've used, 1 Liter and up the liquid condensate colliding with pressurized gas flowing the other direction bottle-necks at the 24/40 ... but a 45/50 works just fine
Using 24/40 is OK for 25 to 1000 ml reactions and most 25 to 500 ml distillations. Purchasing oversize (standard) jointed glassware across the board and as you move up in the volume of a synthesis/distillation/extractions you'll have more versatility with less equipment.
The "Friedrichs" condenser (sometimes incorrectly referred to as Friedrich's condenser or Friedrich condenser), is a spiraled finger condenser invented by Fritz Walter Paul "Friedrichs", who published a design for this type of condenser in 1912. Even if there's twenty and his, it's still a "Friedrichs" condenser, not a"Fred" condenser because his name was "Frierichs", with and "s". The man was brilliant. 106 years later and it's still going strong.
The rotary evaporator system was not invented by Lyman C. Craig until 1950. It was first commercialized by the Swiss company Büchi in 1957, more than 40 years after the Friedrichs condenser became popular (Where do people get this crazy stuff they post?). Aside from any opinion to the contrary, Friedrichs' condenser could NOT have been designed for rotovaps - unless he was psychic.
Unbeatable condensing ability is only one key features of the Friedrichs condenser. It's also comparably COMPACT in size, has REDUCED bulk/mass/weight, simplicity and significantly reduces the height of the apparatus. That 150 cm Allihn I spoke of was four times as large as the Friedrichs that replaced it. I have also seen more restricted and less restricted type necks but no personal experience there.
An Allihn Condenser on a rotovap would run the risk of breaking/create undue stress the bearing junction. Nearly 5 Feet and (Roughly - not certain but HEAVY) 35 lbs of condenser, containing 15 lbs of coolant = roughly 50 lbs, on a slant. I guess one could use clamps, rope, shoelaces or duct tape - thus ruining the elegance of a $15,000.00 Buchi (not to mention the bulbs being on a slant, pooling and flooded, defeating the purpose of a rotovap) - or obtain an 18" high, same weight Fredrichs that contains half the coolant, reduces the Fulcrum by 3 feet and has the best efficiency of any common condenser.
If the interface is sufficient, a quality Friedrichs (W/threaded coolant fittings; threaded release rings or PTFE joint sleeves) is most durable and klutz proof but - as elaborated below - Dimroths are VERY delicate.
Both Dimroth and Friedrichs condensers can be used for Reflux and/or Distillations."Dual use" condensers in a Vertical orientation provide two ultra efficient condensers for one price. If chosen properly, when/if you expand to larger volumes your condenser should be flexible enough to accommodate without adapters.
A few notes on Dimroths. BEWARE: The coolant tubes on Dimroth condensers are necessarily delicate and easily damaged beyond repair. Paper thin coiled tubes effect superior heat transfer. The coil's entire length/mass is supported at the end where it exits/enters the body of the condenser. At the other 'dangling end' of the coil, the amount of flex can exceed the brittle nature of glass where it's connected. Relatively minor lateral acceleration/deceleration impact shock, even to a well padded Dimroth can cause the tubes to break.
When not in use, several thin plastic strips, carefully inserted through both ends will virtually eliminate flexing of the inner tube (Plastic pallet binding strips work great).
NEVER shake, bang on your palm, snap/tweak with your index finger or even tap a Dimroth. They're elegant, precise tools and should be treated gently. Many are broken while washing (That's why the used ones are often so discolored. The vigor required to clean them would risk breakage).
Having never owned a large one, in my unqualified opinion the longer the Dimroth, the more fragile.

That's my personal experience and the opinions I've garnered.

Stacking condensors is structurally dangerous, and definitely useless unless your first condenser is too small to begin with. Sorry for shouting, but that BAD LABORATORY PRACTICE!

150cm? Is that correct?

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How to characterize gaseous ethylene that forms as byproduct?
I'm running a reaction that has ethylene gas as a byproduct. I have observed gas evolution as my reaction proceeded and I have other ways of verifying the identity of the main product, but I thought it might be cool if I could also test for the presence of ethylene.
Are there any simple tests I can use for this purpose? Obviously it is not easy to detect since it's a fairly non-reactive gas, so it doesn't have to have 100% specificity.

gas chromatography

may I recommend GC-MS as being a little bit better than just normal IR-GC for this purpose, but both will work.

GC-MS should work quite well, I agree. It will freeze in liquid nitrogen though, if that makes things easier for you.

It's totally OK to think "how could I identify it", but why would you actually want to do that if you know for sure it is ethylene? Isn't that kind of pointless?

If you want something simple and don't mind less than 100% certainty then you could bubble the gas through Bromine Water - in the presence of an alkene the bromine water would be decolourised from brown when reacting with the ethene.  Of course this does just indicate the presence of any alkene so isn't exactly specific! There is more about this on this website.

You can do NMR experiment on the resulting solution. This should give you at least the negative on ethylene, but maybe also the compound that blows out.

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Ceramics in "bullet proof" vests?
Kevlar is an aramid fiber that is used in ballistics protection vests. From the structure (structure shown from Wikipedia article), you can see why it would work to dissipate the energy from a bullet since there is an extensive network of interactions between the linear chains. The wearer ends up bruised from the force, but unpunctured and avoids serious injury from handguns bullets.

More modern ballistics protection devices use ceramic plates (along with Kevlar) to give greater stab resistance and protection from higher velocity rifle rounds. My question is, what ceramics do they use? Do they have a "flexible" ceramic that will deform to absorb the energy and not break? Or do the plates work by breaking to absorb the energy and, thus, have to be replaced after each event?

As far as I can tell there are two material types used for this - pure ceramics or ceramic mixtures that are designed to break and be replaced, and composite metal + ceramic materials that are designed to be able to handle a couple of rounds without losing too much effectiveness. There are also techniques involving purely metal layers in overlapping structures with fibers of materials like Kevlar or ultra-high-molecular-weight polyethene (UHMWPE) (of which Spectra is a brand).
Shopping around, it seems the best materials are boron and silicon carbides (extremely hard materials, though not quite as hard as diamond), which can be formed into plates with multiple levels of structure by starting with hexagonal plates and forming larger plates from those. In addition to personal armour, the same materials and techniques are used for military vehicle (including aircraft) armour. In the ceramic/metal composite materials, you get ceramic plates formed up within a metal sheet, where a ceramic element will shatter to take most of the impact, while the rest will be absorbed by the metal matrix. The metal matrix then remains intact (but may be deformed) so that the armour is still useful but weakened around the area of impact.
Good places to start from are the Wikipedia pages for Chobham armour (the common name for the tank armour that includes composite metal/ceramic matrix materials), and Small Arms Protective Insert.

Thanks! That makes more sense than the mental picture I had of the ceramic plates and how they would function.

There are a lot of downsides to ceramic armor; among them, they are relatively fragile when dropped, stored in the trunk of a vehicle, or submerged in water. Furthermore, as you point out, they are fail quickly when hit multiple times; steel armor of similar mass can withstand many thousands of hits while the ceramic is finished after a dozen.


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MacGyvering a sealed tube
Is it possible (advisable?) to make a sealed tube by melting and twisting the end of a borosilicate test tube? One issue that I can think of is that the pressure might be too high if I then proceed to heat the tube. I thought that I could maybe connect the tube to a vacuum tube and heat it in the middle. Any experience with this?

I don't have any experience doing this, but I understand it's a fairly common procedure for air-sensitive NMR or X-ray diffraction samples. The key phrase you probably want for searching for techniques/protocols is "flame-sealing".

Usual way to work with air sensitive compounds is to work on Schlenk-line (search Wiki). When you transfer your substance you can leave it in Schlenk tube, or if it's really sensitive and you need to keep out of the line for longer period of time you have in general three possibilities: (1) use glassware with J. Young valve (available even for NMR tubes); (2) store your standard Schlenk glass in glove box filled with inert gas and (3) the oldest is to melt your glass.
Now back to your question. Usual way is to prepare glass with narrow place on it prior to doing experiment (that's were later you will melt it). By narrow I mean the hole inside should be ~2-3 mm (thinner glass, smaller hole). I used to use Schlenk lines, so on opened end I have standard 14 female cone that was connected to a vacuum line. I used to set vacuum ~0.1 Torr and with propane/butane/oxygen torch one can easily close it. The huge advantage of this old way is that you can prepare your glass even with 3mm thick walls that allows you to work with substances with very low boiling points (in this case to close the glass one need to freeze the stuff with liquid N2 or if not available at least cool it with dry ice bath). Of course later opening is quite tricky and usually also had to be done after freezing of chemicals.
Last part if it's advisable. That one is as usually hard to answer. I have no idea how experienced you are. Judging by question not to much (sorry). So I wouldn't start with anything to reactive in the beginning. Generally try to play first with glass and torch and get a bit used to it. Do it on proper chemical glass (not just bottle glass).
If you are "advanced guy" and you want to make yourself f.e. PfBF2 borane you have to use this technique - any other gives just a miserable yield.

Possible? Doubtless. Advisable? Depends what you're sealing, I'm sure. The pressure won't be a problem though, if anything it'll be a partial vacuum inside and you'd have to heat it near the tube's melting point to get it back to atmospheric pressure.  Consider that as you're heating the tube to seal it, you'll heat a good portion of the air inside as well, and once that cools it will be well below 1atm, only getting back there after heating to the melting point, the temperature at which it was sealed in.

It is indeed common, though you don't use standard borosilicate glass as I understand it. Also, it is a technique that needs special training to do right; it can be dangerous to both you and your compound if you don't get said training first. Ask around your local inorganic chem department, I'm sure someone can teach you to do it.


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Do Na or Li need heat in order to react with Cl₂?
I've read how putting a burning piece of sodium in chlorine would produce white $\ce{NaCl}$ fumes. However, I can't find out whether Na or Li can react with $\ce{Cl2}$ without heating or burning.

From personal experience, Na(s) and Cl2(g) react at room temperature without added energy. Li(s) and Cl2(g) should do so also.
Consider the following heat of formation data harvested from the NIST Chemistry Webbook:
For NaCl:
$$\ce{Na(s) + } \frac 1 2\ce{Cl2(g)->NaCl(s)}\quad \Delta H_f^o=-411.12 \; kJ/mol $$
For LiCl:
$$\ce{Li(s) + } \frac 1 2\ce{Cl2(g)->LiCl(s)}\quad \Delta H_f^o=-408.27 \; kJ/mol $$
These data suggest that both reactions should occur without added heat. However, the rate of reaction is not determinable from these data. The reaction between Li(s) and Cl2(g) will likely be slower.

In general, the NIST Chemistry Webbook is a great resource for physical constants, phase change and thermochemistry data, and spectroscopy (IR, MS, but not NMR).

Presumably the thickness of the surface hydroxide/carbonate coating that forms in air will affect whether any particular metal sample will react at room temperature.

These data suggest nothing. Thermodynamically, Mg or Al should react with oxygen without added heat, yet in fact they would just sit there for ages, not doing anything.


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Synthesis of chloroquine from 4-amino-7-chloroquinoline
Can we get chloroquine from 4-amino-7-chloroquinoline (7-chloroquinolin-4-amine) by adding 4-chloro-N,N-diethylpentan-1-amine $(\ce{CH3CHCl(CH2)3N(Et)2})$ ?

Looking at the formulas for your two compounds, it seems like a simple substitution. What are you looking for in particular?

I've added a drawing of the proposed synthesis with hydrogen chloride as the inferred elimination product. I flipped the pendant alkylamine to save space on the page. ChemFig rocks, btw.

My intuition tells me that the nitrogen in the ring is more nucleophilic than the exocyclic amine. A quick-and-dirty electrostatic potential calculation using B88LYP/6-31G(d) on the PM3 geometry shows a large electron density (blue and violet) on the ring nitrogen and lower density (yellow) on the exocyclic amine:

Are there conditions that might work? Yes. Should you try it, if you need that molecule? Yes. Will it be difficult? Yes.

It's a pretty untypical approach to chloroquine and you didn't tell how you would synthesize 7-chloro-3-aminoquinoline and the secondary chloride.
There is a good chance that the planned reaction will furnish a quarternary ammonium salt as a side product:

The reaction will not necessary stop at this point.
Ia addition, there's a good chance to bis-alkylate the $\ce{-NH2}$ group of the quinoline.
All this is not very promising and any attempt to optimize this reaction is fighting problems you would not have if you had decided for a more reasonable and well-established alternative in the first!
That would be the synthesis of chloroquine from 4,7-dichloroquinoline and 4-diethylamino-1-methylbutylamine, where both starting materials are available through well-known procedures.

Yeah, and even synthesis of aliphatic amine chloride looks like a real problem

@Mithoron Yes, 4-diethylamino-1-methylbutylamine is definitely easier to synthesize.

As an additional note, the aliphatic chloride could also be displaced by nitrogen in an intramolecular cyclisation to give a five-membered ring. Not to forget that multiple alkylations of the desired product would be observed, too.

@Jan Yes, OP is looking for as much trouble as possible :D

yes, it is quite possible as aromatic anines nitrogen being moore nucleophilic further the secondary chloro reacts easily, the reaction may either be SN1 or SN2. I have done couple of reaction of this type, if you need help feel free to contact me.

Could you elaborate on how this reaction works, maybe with a breakdown of the mechanism?

I would be skeptical. I would assume multiple alkylations to take place even aside of the problems others pointed out.


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Bond energies database
I’m looking for a public database of bond energies of biochemical compounds. Does one exist?
Otherwise, which is the standard computational methodology to calculate them? Any references would be very appreciated.

Hi, and welcome to Chemistry StackExchange! What are you looking for specifically? An table of covalent bond energies? These exist (and are widely available in textbooks and on the net) but are not specific to biomolecules. Or are you talking about binding energies of pairs of molecules?

thanks, precisely covalent bond energies is what I'm looking for. Despite I'm interested in biochemical compounds, please let me know the net resources that you comment.

Have you searched NIST?

The Comprehensive Handbook of Chemical Bond Energies might be able to help.  It walks through bonds in different functional groups.  It can be found at http://www.ebah.com.br/content/ABAAABH2cAD/comprehensive-handbook-of-chemical-bond-energies-0849373662


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How would I calculate the Normality of a 0.00167 molar solution of KIO₃?
I need a 0.01 N solution of $\ce{KIO3}$ to use as a standard in a Winkler titration.  The protocol that I am using specifies that this is 0.3567 g L-1, which I have calculated to be 0.00167 M based on 214 g mol-1 for $\ce{KIO3}$.
I like to understand the calculations in the protocols that I am using, so how would I calculate that a 0.00167 M solution of $\ce{KIO3}$ is 0.01 N?
EDIT (More detail on the reactions)
The $\ce{KIO3}$ solution is added to an $\ce{NaOH} \cdot \ce{NaI}$ solution under acidic conditions for:  $\ce{IO3- + 5I- + 6H+ \rightarrow 3I2 + 3H2O}$
The $\ce{I2}$ is then titrated to standardize a thiosulfate solution.

Do you know what the iodine in $\ce{KIO3}$  is becoming? In all redox reactions, the final oxidation states are important. Looks like it can become a mix of iodide and iodine, which makes this problem slightly more complex.

In this case, the reaction was particularly important since it's a Comproportionation reaction, which is wierder than usual. I'll write up an answer if I get time.

Hmm, are you sure that it is 0.00167 M ? I'm getting 0.0002 M. I may be wrong, though

I calculated it as 0.3567 g $\ce{KIO3}$ / L * mol \ 214 g $\ce{KIO3}$. So 0.3567/214=0.00167.  Thanks for your help with this.

One source with details such as given by OP is here https://ocw.mit.edu/courses/earth-atmospheric-and-planetary-sciences/12-097-chemical-investigations-of-boston-harbor-january-iap-2006/labs/dissolved_oxygen.pdf

Hmm.. I think that your data is wrong, I get that the normality:molarity ratio must be 5 in this case.
This ratio (I call it the $\eta$-factor) for a compound participating in a reaction is the net change of its oxidation state.
In this case, one iodine in $\ce{IO3-}$ (+5) becomes one iodine in $\ce{I2}$ (0). The net change in oxidation number is 5, so $\rm normality=5\times molarity$. Normality is like the reducing/oxidising "power" of a solution--or the concentration of electrons that it can furbish/absorb.
Your density, your $\mathrm{g~L^{-1}}$, or your value for normality is wrong here.
Note that one needs to be careful with this reaction. It is comproportionation, where the same element is oxidising itself (Basically, two forms of the element in different oxidation states are reacting with each other). We can have some trouble when dealing with the products of such a reaction, though dealing with the reactants is fairly simple, as shown above. I'm writing a detailed self-answered post on this later.


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CHARMM parameter files for NBD (nitrobenzoxadiazole)
I have been looking all over the internet for forcefield parameters for a fluorescent probe called NBD (nitrobenzoxadiazole). Here is an example of an NBD labeled fatty.
I have so far only found parameters for AMBER or GROMACS. Also when following the articles from these two reviews, there seem to be no CHARMM parameters anywhere.
My questions are:

What CHARMM forcefield parameters would you use for NBD?
How could I transfer the AMBER or GROMACS parameters for use in CHARMM?

The links to products.invitrogen.com and springerlink.com are both broken. Perhaps you could take a look, whenever possible...

Converting between force fields is always an iffy affair, as different force fields were parametrized with different reference sets, and possibly even with different intent in mind.
You have two options here:

You can extract the raw parameters from the particular GROMACS or AMBER force field used here, and try to translate them into the parameter formalism CHARMM uses. Generally speaking, force fields consist of a set of parameters for each atom type, and a set of equations these parameters are used in to calculate forces and geometries. Those equations represent aspects of the atom's or molecule's properties that have an influence on the calculated forces or the molecule geometry, like bonding interactions, non-bonding interactions, angle contraints, bond length constraints, and so on. You need to find the corresponding equations in CHARMM's force field framework, and then translate the parameters from AMBER or GROMACS so that the force field is the same in CHARMM.
Alternatively, you can look at the various CHARMM force fields and try to find one that works with NBD and any other molecules you got in your simulation. The first article you cited details where the various general parameters they used were from and how their simulation in GROMACS was set up. You should be able to define the same simulation parameters in CHARMM and rerun their simulation there. You can then use their published data to benchark your own simulation against theirs.


text:
Liquid which neither dissolves in water nor oil
I am looking for a liquid which is not miscible (or very close) with water or vegetable oil.  Ideally, I want to have a container which shows three layers (the three liquids) with clear boundary layers.
I was thinking that I should find a non-polar compound of very high molecular weight - such that difference in the density of that compound and oil is enough to draw it down out of the oil.  However, I haven't found it yet.
I would prefer a compound which is non-toxic and non-flammable if possible.
It feels like such a simple question, I must be missing something.

One thing you might want to add to your question is how much stability you want/need - for example, if you're careful, you can get quite a lot of layers set up, but if you plan to shake it, most of these will combine.

Remember, miscibility isn't that much to do with density as it has to do with polar/nonpolar. If you get a third (polar) layer above the oil, on shaking it'll mot probably dissolve in water.

@Aesin: You can make that into a nice answer, especially if you explain how miscibility works and why shaking it kills the system.

Sorry, I should have been more specific.  There will definitely be moderate mixing in the system.  Using metal is possible, but I would like there to be a fairly significant amount of stuff (aka 5 - 10 gallons worth) because I'd like it to be a easy to see visual display, so cost may be an issue there.

David E. H. Jones demonstrated a jar with seven or eight liquid phases in the 70ties. (of course shaking-proof!)

Fluorocarbons are compounds that are hydrophobe and lipophobe. It's a special property of perfluorinated compounds: They are non-polar and thus hydrophobic but in addition as the high electronegativity of fluorine reduces the polarizability of the atom, fluorocarbons are only weakly susceptible to the fleeting dipoles that form the basis of the London dispersion force. As a result, fluorocarbons have low intermolecular attractive forces and are lipophobic.
I'm not sure about their toxicity, but this seems to be an issue, e.g. here.

Good suggestion.  Many are highly toxic and not desirable, but some of them are used in medical applications and would be acceptable.

@Greg: you'll probably be interested in those perfluorocarbons used as "blood substitutes", due to their capacity for retaining oxygen. They are nontoxic, AFAICT.

How about Mercury and Gallium? Mercury is rather toxic, but you can buy Gallium online and it's a liquid at 85 degrees Fahrenheit - so it might work if you left it in some sunlight when you wanted the Gallium to be liquid.
If you don't plan on opening it often or need to worry about it breaking, I'd try Mercury.

Elemental mercury is liposoluble.

Interesting - though it would still form layers if you added the Mercury, then water, then oil. Just don't shake it around too much I guess.

see Aesin's link in his comment below the question for adding many layers if you don't shake things. Alternate layers of water and oil do the trick, too, if you pour them slowly and do not shake.

Would the mercury/gallium not pull out of the oil over time due to their weight, and eventually resettle? Or would the mercury be bound into the oil? If that's the case, you could just add so much mercury that no more can dissolve in the oil.


text:
Basic chemistry equation with ammonia
A table in the book 'Chemistry NCEA Level 1' has 4 headings, the first of which is 'Volume (mL) ammonia solution provided (conc 2 g L-1)'
Could someone please provide a comprehensive explanation of what the end part in brackets means.
This is for a chemistry newbie, as you may gather from the book title.
(this isn't really homework but I couldn't find a relevant tag. It's just someone studying chemistry for fun.)

Hi Mark, welcome to the site.  I think that this question is too localized, as it would only apply to people using that book in the future.  To answer the question, though, it's saying that the concentration of the ammonia solution is 2 grams per liter, so there are 2g of solute (ammonia) in 1 liter of solution (ammonia + water).

I don't understand why this has been closed. If someone without physics knowledge asked the meaning of e=mc2 would that be closed? If someone didn't know that 3n meant 3 times n would their question be closed? I don't know what 'conc 2 g L-1' means. Does the L-1 part mean L to the power of -1? Why does it not mean 2 times g times L-1?

I understand where you are coming from, but the answer to your first two questions is almost always, yes.  These issues are known as "general reference" questions and can be looked up online.  You are welcome to start a Meta discussion about  why this question was closed, if you'd like.  The "L-1" is an old typewriter notation that did indeed mean "L raised to the -1 power" or 1/L.  If you peek ahead in your text to the section on concentration, it should answer that question and more.

@Mark: The L-1 is just a way to easily note  the unit g/L (grams per liter). Regarding the closing of the question, we allow simple questions here. They just shouldn't be easily Googled. One major point to take into account while asking the question is this: "Will it be useful to future visitors/will it help them learn?". We are not one of those quick help fora, actually, we aren't a forum at all. We aim to make the Internet better by keeping good, conceptual questions that will be useful to others. Simple questions are OK as well, though they shouldn't be easily Googled.

This question serves as a bad example for the homework policy.


text:
Synthesis of mandelic acid from benzaldehyde
For an upcoming exam, I have to describe the synthesis of 2-hydroxy-2-phenylacetic acid (mandelic acid) from benzaldehyde. The provided answers describe quite a long process, consisting of 8 stages. However, I've been suggested a much simpler method, consisting of:

$\ce{C6H5CHO + HCN -> C6H5CH(OH)CN}$
$\ce{C6H5CH(OH)CN + HCl + 2H2O -> C6H5CH(OH)COOH + NH4Cl}$

This is even simpler than the 3-stage process described in the Wikipedia article. Could it be a viable answer?

The direct reaction with HCN seems to be how the mandelic acid was synthesized in the 19th century. If you check how the reaction with sodium bisulfite is actually performed (Organic Syntheses, Coll. Vol. 1, p.336 (1941); Vol. 6, p.58 (1926).) you'll see that this is not more complicated than the direct HCN route as those first two steps are performed in one reaction step.
From this preparation, several preparations via the HCN route are cited:

Naquet and Louguinine, Ann. 139, 299 (1866); Wallach, Ann. 193, 38
(1878); Spiegel, Ber. 14, 239 (1881); Ultee, Rec. trav. chim. 28, 254
(1909); Rule, J. Chem. Soc. 113, 12 (1918).

I looked at the first page of the oldest of those papers from Naquet and Louguinine from 1866 and it states:

Bekanntlich bildet sich diese Säure bei der wechselseitigen Einwirkung des Benzoesäure-Aldehyds und der Cyanwasserstoffsäure bei Anwesenheit von verdünnter Chlorwasserstoffsäure gemäß der Gleichung:
$$\ce{C7H6O + CNH + 2 H2O + HCl -> NH4Cl + C8H8O3}$$

Translated, this means that you can synthesize mandelic acid from benzaldehyde, HCN and HCl.
I'm assuming you're not planning to actually synthesize it using this route yourself, I would guess that avoiding the use of HCN directly was one of the motivations behind the synthesis route using sodium bisulfite.
$%edit$


text:
Any false chemical element in history of chemistry?
Was there any false chemical element introduced in history of chemistry? I mean a substance that after a while proved that it was not an element.

It wasn't considered an element, but let us not forget phlogiston.

Also, not to mention electric fluid

There's an entire list here: http://www.vanderkrogt.net/elements/didnot.php

Caloric (heat) was listed as an element by Lavoisier.

The example that comes to mind is didymium, which turned out to be a mixture of praesodymium and neodymium. The term is still used, as far as I know, to refer to glass doped with a mixture of these lanthanides. This is discussed in one of the 'Chemistry in its element' podcast episodes.
As it turns out, wikipedia has a category called 'Misidentified Chemical Elements'. Especially notable are illmenium, dianium and pelopium, which were all likely mixtures of niobium and tantalum, two rather similar elements.

One of the most important examples of a misidentified element is  coronium.  This element (or what was thought to be a new element) came to play a major role in our understanding of the Sun.  Originally identified by spectroscopic analysis of the Sun's corona during a solar eclipse, coronium turned out to be iron so heavily ionized that to account for it, the corona must be heated to millions of degrees.

The question is: Was there any false chemical element introduced in history of chemistry?

The question is very broad, but the answers given so far are mainly based on modern time suggestions. For example, coronium was suggested as an element that 19th century scientists believed might exist in the corona of the sun, hence the name. According to Science Notes:

Coronium was suggested as a new element as the source of an unidentifiable green line in the spectra of the solar corona, which is visible during a total solar eclipse. This green emission line at $\pu{530.3 nm}$ was first observed by two scientists individually during the solar eclipse in August 7, 1869. This line did not correspond to the emission spectra of any known element at the time. Dmitri Mendeleev, known for developing the periodic table, proposed the name newtonium rather than coronium for this element who believed newtonium was an element slightly lighter than hydrogen.
It wasn’t until the 1930s that  Walter Grotrian and Bengt Edlén determined the mysterious green spectral line observed during solar eclipses was due not to some unknown element, but to highly charged iron, $\ce{Fe^{13+}}$. Highly charged metal ions are responsible for other previously unidentified lines in the solar spectra, too.

However, I believe that the most important mis-identification of elements goes back to BCE time. Based on his previous philosophers ideas, Plato (428-347 BCE) has suggested that the terrestrial realm was imperfect, which was based on the four elements: (1) Earth; (2) Water; (3) Air; and (4) Fire. This theory was perfected by his best student, Aristotle adding some essence (Ref.1):

An element, we take it, is a body in which other bodies may be analyzed, present in them potentially or in actually (which of these, is still disputable), and not itself divisible into bodies different in form.

References:

Andrew Ede, The Chemical Element: A Historical Perspective; Greenwood Press: West Point, CT, 2006, pp. 13-15. ISBN 0-313-33304-1.

license: Creative Commons - Attribution Share-Alike - https://creativecommons.org/licenses/by-sa/4.0/


text:
Is there any danger of forming explosive reaction products during the electrodeposition in Diethylene glycol?
I know that perchloric acid could be potentially explosive and some organic solvents could form dangerous peroxides.
I am going to do an electrodeposition in diethyleneglycol solution containing $\ce{NH4Cl}$, $\ce{ZnCl2}$ and $\ce{FeCl2}$. Since in my case there are all elements to form $\ce{HClO4}$ and I am going to heat everything up to 60 degrees I want to be sure that it is safe.
Is there any danger of forming explosive reaction products during the electrodeposition?

I am way too lazy to do it myself, sadly... but this will help you.

Diethylne glycol reacts readily with perchloricacid, however the oxidation and hazard potential of perchloric acid depends on its concentration in side the solvent.
If concentrations are high enough and more than 10 percent perchloric acid formation is expected with 100 percent reaction yield, then it can be explosive.
What actually happens that as soon as diethylene gylcol meets with perchloric acid, diethylene glycol decomposes and some dangerous peroxides formed. These are very reactive to water, so if you can provide very dry and inert atmosphere it should be safe but you should give extra care to make sure that you do not have humidity and water both in your chemicals and in the atmosphere.
If you are going to do the experiment and use milimolar range chemicals then just do it, because you will not be able to form enough perchloricacid to get problems.


text:
How to translate the mobile phase from thin layer chromatography to column?
I recently was told that once you find a good solvent mixture for separation on thin layer chromatography (TLC), you should only use 10% of the more polar solvent (for silica) when you run your column. I tried it and it worked great!
Since I'd never heard that before, is that a standard rule of thumb, or did I just get lucky?

There is an article in the Journal of Organic Chemistry that describes this process. I'm having trouble finding it. When I do, I'll make an answer out of it.

Is this the article you are referring to? Still, W. C.; Kahn, M.; Mitra, A. J. Org. Chem. 1978, 43, 2923–2925.

That is indeed the article.

I didn't see anything in there about changes between TLC solvent and column solvent: I thought you used the same solvent in the TLC and column, and that seems to be what the article describes.

@Canageek, yeah, they seem to use the same solvent mixture. I guess that was a fluke that that tip worked.

Turns out I was wrong, and moving to a less-polar solvent is standard procedure. We don't use the 10% rule though; We typically use 50:50 DCM:Hexans for the TLC and 1:3 DCM:Hexans on the column.

I never heard about the "10% rule" and I never applied it. Transferring TLC (silica on aluminium plates, Merck) results to preparative column chromatography on similar silica usually worked without changes, although the situation is a bit different in the column: the upper part of the TLC plate is rather dry when letting it run, while the silica in the column is not. Note that I was never a friend of the dry packing method and always prefered the slurry method to prepare the columns.
Anyway, I never thought about increasing the polarity.
Bad separation typically was a result of column overload, bad packing or cracking of the stationary phase, e.g., when the diethylether evaporates in the column.
+1 For citing the classic article on flash chromatography!
I can really recommend the technique and I used it ... more than once ;)

I believe the 10% rule only applies to ''strong'' solvents such as methanol or ethanol.  These solvents have polarities around 6.6 compared to ethyl acetate which has a polarity of 4.3 or so and hexanes which is what, zero?
Anyways the rumor is that strong solvents will degrade the silica.  Here is an interesting article that confirms and denies said rule. http://www.flash-purification.com/does-methanol-really-dissolve-silica-during-flash-column-chromatography/
As with all good chemistry questions it depends and nothing beats running the experiment.  The proof is in the NMR.


text:
What is the difference between an ionic and molecular activity model? (Example: Diopside-Hedenbergite)
Original question:
When looking at a Diopside-Hedenbergite solid solution, what is the difference between an ionic and molecular activity model? Why does the moleclar model yield higher values?
Post-Edit question:
The two activities calculated below seem to me like I'm just comparing some random values. What is the true activity of the Fe- or Mg-mineral?
So far I got:
Dipside and Hedenbergite are Pyroxenes. Pyroxenes have to metal sites (M1 and M2):
\begin{align}
M^1M^2[Si_2O_6]
\end{align}
The difference between Hedenbergite and Diopside is:
\begin{align}
Hedenbergit (Fe^{[M1]}Ca^{[M2]}[Si_2O_6]) \longleftrightarrow Diopside (Mg^{[M1]}Ca^{[M1]}[Si_2O_6])
\end{align}
The activity of an element is assumed to be equal to the concentration:
\begin{align}
a=\color{Red}{\gamma}x \\
a=x
\end{align}
In the next step one has to calculate the concentrations on the M1 site because it's the site where the ion exchange happens:
\begin{align}
x[Fe]=\frac{[Fe]}{[Fe]+[Mg]}
\end{align}
\begin{align}
x[Mg]=\frac{[Mg]}{[Fe]+[Mg]}
\end{align}
The molecular mixing model apparently compares the activity of the Fe- and Mg-mineral on the M1 site:
\begin{align}
a[Mg]=x[Mg]
\end{align}
\begin{align}
a[Fe]=\frac{1-x[Mg]}{x[Fe]}
\end{align}
Whereas the ionic mixing model compares the activities of all the cations of the minerals:
\begin{align}
a[Mg]=x[Mg]^{[M1]}x[Ca]^{[M2]}
\end{align}
\begin{align}
a[Fe]=x[Fe]^{[M1]}x[Ca]^{[M2]}
\end{align}
Why are there two activities? Shouldn't there be one true activity? What does the one activity mean and what does the other mean?

Your question intrigues me, so I made some research. This is the only verifiable information I've found so far: http://www.doitpoms.ac.uk/tlplib/solid-solutions/index.php

Looks OK to me (not sure though, not too familiar with this area). Note that "no answer" does not mean off topic-- just that the right people may not have seen it. Your error propagation one is probably off topic though, so Ill leave it deleted for now. Next time, flag the post, though its good you are willing to clean up your own posts :)

The activity of a chemical species is a key concept in thermodynamics of solutions. While the thermodynamic potential is the “true thermodynamic variable”, i.e. the one that plays a primary role in statistical mechanics, chemists do tend to like activity (which is univocally related to chemical potential) because it is close to another quantity which has clear physical meaning: concentration.
Now, relating the activity (or chemical potential) of a species to its concentration in a mixture is impossible to do in a general way. Models have been developed to provide approximations, however, the best known being the ideal solution, where activity equals molarity (concentration). Your “molecular mixing model” appears to be a model of ideal solid solution. However, one cannot expect this approximation to hold in case of strongly non-ideal solutions, and ionic solutions are prominent in that category. The “ionic mixing model” you describe thus appears to be another such approximation, more suited to these ions in a solid solution. I have never encountered it before, but I found a series of lecture slides where it seems to be very neatly introduced (see around slide 31).
While I cannot fully address the question, I hope this helps!


text:
MO-Scheme of SCN- and its bonding properties when used as a ligand
I tried to figure out the MO-scheme of the tetragonal-bipyramidal complex trans-$\ce{[Co(en)2(NCS)2]SCN}$ in which the isothiocyanate ligands are bound to the $\ce{Co^3+}$-Ion in $\eta^{1}$-mode (en = ethylenediamine). While undertaking this task I noticed that I don't know the MO-scheme of the isothiocyanate ion, which I need to determine whether it acts as a $\sigma$ and/or $\pi$ acceptor or donor.
I tried to construct the MO via a symmetry-based treatment but ran into the problem that the isothiocyanate ion - belonging to thet $C_{\infty v}$ point group - lacks symmetry elements perpendicular to its principal axis. Thus, I wasn't able to construct the bonding and antibonding sigma-orbitals with this method.
I could construct those $\sigma$-interactions by hand or use $\ce{CO2}$ as a model but I don't know whether this will give me the right results. I haven't found the MO-scheme in the internet so I would be happy if someone could advise me on how to construct a qualitative MO-scheme of $\ce{SCN-}$ or simply show it to me. I would also be very grateful for getting some information on how exactly the isothiocyanate ion acts as a ligand (which $\pi$-interactions and how strong are they, is there also $\sigma$-donation as with carbonyl ligands, etc.).
Edit:
I forgot to mention explicitly that the $\ce{SCN-}$ ion is bound to the metal center via its nitrogen atom.

One issue with SCN-, is that it can bind either at the nitrogen or at the sulfur. The way you presented the formula, $\ce{[Co(en)2(NCS)2]SCN}$, suggested that the SCN- ligand is bound via the nitrogen. If it were bound by the sulfur, the formula might be written  $\ce{[Co(en)2(SCN)2]SCN}$.
SCN- should have a similar MO diagram to OCN-, which is isoelectronic with CO2, so your idea to compare to CO2 is not  a bad one. However, since SCN- has a different symmetry group, you cannot trust that comparison too much. Your comparison to CO is also a good one, since SCN- and CO have the same point group.
However, what you really need to do is determine the symmetry group for $\ce{[Co(en)2(NCS)2]SCN}$. Then, assign symmetry labels to the SCN- ligands. The MO that is most important for SCN- is, qualitatively, the "lone pair" on the atom pointing at the metal center, which can form a "$\sigma$-bond" with cobalt. I would only consider using those $\pi$ orbitals on SCN- if the complex needs more electrons to get to 18. When you assign symmetry labels to the SCN- ligands, consider them as a pair: i.e. assign a symmetry label to the in-phase pair and then assign a symmetry label to the out-of-phase pair. The in-phase pair should probably have the same symmetry label as the cobalt $d_{z^2}$ orbital.

Do you really think that the most important interaction of $\ce{SCN-}$ ligand with the metal center is a $\sigma$ interaction? I would expect that the HOMO and (maybe) also the LUMO of $\ce{SCN-}$ is comprised of $\pi$-MOs (non-bonding (-> lone pairs) for the HOMO and antibonding for the LUMO). The $\sigma$-interaction should only come about by an additional interaction of a $\sigma$-Orbital with the metal $d_{z^2}$. But since the $\sigma$-MOs should be rather low in energy I wouldn't expect the interaction to very strong. Am I missing some points or do you think my arguments are reasonable?

The HOMO of SCN- is probably a nonbonding orbital, ie one of the "lone pairs".

Sorry if it wasn't clear from my first comment, that I totally agree with you there: If $\ce{SCN-}$ is similar to $\ce{CO2}$ then I would expect the HOMO to be a pair of nonbonding $\pi$ orbitals with out-of-phase "p-orbitals" on $\ce{S}$ and $\ce{N}$ and a nodal plane through $\ce{C}$. Would this mean that $\ce{SCN-}$ will be mostly a $\pi$-donor ligand (maybe with some additional $\sigma$-donation from lower lying $\sigma$-MOs) or will the higher lying $\pi^{*}$-MOs (= LUMO) be more important for the interaction with the metal center so that $\ce{SCN-}$ is more of a $\pi$-acceptor ligand?

@Philipp Inspection of the spectrochemical series on Wikipedia seems to suggest that the thiocyanate ion seems to be more of a pi donor/sigma donor ligand since the ligand seems to be placed together with other such ligands. However, it is interesting to note that the pi donor character increases when coordination is done through the N atom while the sigma donor character increases when the coordination is done through the S atom.

@Philipp I have misread the Wiki page regarding the last part of the previous comment. The pi donor character is higher when coordination is through S and is lower when the coordination is through N. This is easily rationalised by the fact that the lone pair orbitals on S are larger and more diffuse allowing for easier overlap with the metal d orbitals.


text:
Handling and storing sodium hydroxide, hydrochloric acid, hydrogen peroxide and isopropyl alcohol
I'm an electrical engineering and I hope you can give me some advice as I'm not much into chemistry.
I'm trying to start making my PCB (Printed Circuit Board) at home but I'd like to have some advice concerning how to handle and where to store these (potentially dangerous) chemicals.
I need NaOH as a developer, hydrochloric acid + hydrogen peroxide as etchant (they should be a better alternative than Ferric Chloride) and isopropyl alcohol (to make soldering flux).
As far as I understood I need:

Glasses (to protect against vapors and splashes)
Gloves (to protect when mixing the chemicals)
Some kind of protective clothes (not sure if I need to, but obviously I'm not wearing nice clothes when doing etching)

I have a few questions though:

Do I need kind of mask / filter (to put on my nose) - are these vapors dangerous ?
I read gloves should be stored in quite cold locations and not be exposed to sunlight. Does that apply to the chemicals listed above too ? I'd like to store them outside but maybe it's not safe ...
Hydrochloric acid comes in 36% concentration but I think I'll need to dilute it to 1%-5% max. After it has been diluted it should be quite safe, isn't it ?
Is it safe to store them inside a house (as long as it's well ventilated) ?

You probably don't need a mask - vapours from hydrochloric acid can be a little noxious, but if you're diluting it down and using it in a reasonably ventilated area, you shouldn't get much of that. Isopropyl alcohol vapour is toxic like most alcohols, but I don't think an ordinary mask will prevent that, and you'll be better off just making sure you have enough ventilation.
Standard white cotton labcoats are inexpensive and can be quickly removed if you spill a nontrivial quantity of something on them.
Hydrogen peroxide will probably come in an opaque container, but if not, it does break down in sunlight. Everything else is pretty stable to light, although you want to avoid storing the alcohol somewhere that's going to get too warm, because it evaporates readily and can pressurise the container more than is a good idea.
All these materials are pretty safe to store indoors in unspecialised storage, although for commercial purposes or quantities you may need to take regulatory precautions that will obviously vary by locale.
Edited to add: The copper chloride that comes out should be regenerated as many times as possible, but when you eventually end up with too much of it, you'll want to already have thought about what you're going to do with it - it's pretty toxic and you can't (or, well, shouldn't) just wash it down the sink.

Thank you for your reply. For highly diluted chemicals (lots of water and a bit of chemical) there is almost no problem then. However - the non-diluted chemicals can be stored - with their container tightly closed - inside a closet or in a cellar? And what about the after-etching products (I know I should put them in a box similar to those of these chemicals and dispose of them properly, but do I need to keep this inside as well ?). Thank you.

Furthermore does any of these chemicals have a sort of "Expire date" after which they'll be no more useful or something ? I ask this because I'd get a quite big supply if I can. Small supplies costs a lot more per liter ...

Most substances have an expiration date printed on the container. You can also inquire about that when purchasing.

Also note that you should NEVER store strong acid and strong base next to one another. They create a lot of heat when they mix. If you are storing them on the litre scale you should get a proper fire resistant chemical storage cabinate, one for each acid and base.

@Canageek: even if they're in their container they cannot be stored one next to the other? What kind of fire resistant cabinet for instance (any recommendations)?

Um, I don't know what they are called. We have them in all labs. Yeah, in case there is a leak, you don't want the two mixing, or they might cause a fire.

@Canageek: but this container should be non-metal fire resistant I suppose. Or is flame retardant good as well (or is it the same thing?) ?. Sorry but I only found small flame retardant boxes at my preferred shop.

The decomposition of hydrogen peroxide can be retarded by refrigeration. What concentration is it?

Hydrochloric Acid 36.2%, Hydrogen Peroxide 35% (there is also at 1-3% but it costs way too much for what it is).

I've got a picture of the containers, I'll post it up when I get home.

It's worth pointing out that mixing hydrochloric acid and peroxides could lead to the formation of chlorine gas. If you dilute the two solutions before mixing them, the risk of toxic exposure is lower. I am not familar with the use of etchants for electronics, but consider using pre-diluted sulfuric acid, because it basically cannot form dangerous gases with peroxides.

A few additions:
shelf life

$\ce{H2O2}$ won't last forever, not even in the dark and cool
edit: if it is stored cool and dark, and does not get contaminated, it will last for some years. However, the degradation is catalyzed by metal ions (famous: $\ce{Mn^2+}$) and bases ($\ce{OH-}$), so the problem is that once it gets contaminated (there is stabilizer added, but that of course won't last forever and cannot prevent every type of degradation), it will degrade rather fast.
the isopropanol must be kept tightly closed to prevent evaporation (the container it comes in is usually good),
but $\ce{NaOH}$ (pellets in tightly closed plastic container) and $\ce{HCl}$ (glass bottle with plastic cap) can be kept for years without problems.

Storage of NaOH

$\ce{NaOH}$ is deliquescent (very much hygroscopic).
$\ce{NaOH}$ also reacts with $\ce{CO2}$ from air to $\ce{NaHCO3}$ and $\ce{Na2CO3}$. The corresponding reduction in base equivalents is not going to hurt you much, but
the carbonate can form a crust at the lid of $\ce{NaOH}$ bottles. You should not store $\ce{NaOH}$ in glass stoppered glass bottles as opening them becomes very hard and finally impossible with the carbonate crust.
($\ce{NaOH}$ is also directly corrosive for glass, but not much)
Overall, it is best to store the $\ce{NaOH}$ pellets in plastic bottles (they usually come stored this way) and prepare frequently small amounts of solution.
You can also store the $\ce{NaOH}$ solution for a while in a tightly closed plastic bottle.
How frequently you need to make fresh $\ce{NaOH}$ depends: If you use the $\ce{NaOH}$ in an open bath daily, it won't last as long as if you just need it every few weeks, and meanwhile it is stored away from fresh air. I believe your application is rather robust wrt. the real $\ce{NaOH}$ concentration, so you can probably reuse the $\ce{NaOH}$. On the other hand, preparing the solution is really fast & easy: take the water, count the needed number of $\ce{NaOH}$ pellets into it and let them dissolve.

So the biggest problem would be hydrogen peroxyde. Any idea how long it should last about (like 1 year, or 1 month), as an order of magnitude ? And best pratices for storing HCl and NaOH in solution ? I don't think I can put them on my shelf in the living room though :D.

More accurate would be "NaOH is deliquescent" (so you'd better keep it in a watertight box or you'll be left with slush)

@Manishearth: thx, new English word learnt :-). Watertight usually isn't a problem if it has to be air tight, too.

Reactive chemicals should be stored in chemical storage cupboards to prevent fires. I've attached an example of the type of cupboard we use in my lab to give you an idea of what I am talking about. We use an under-fumehood model, but freestanding models are avalible. Acids and bases should not be stored together. MSDS sheets from the manufacture may specificity additional constraints on which chemicals can be stored alongside one another. Note that hydrogen peroxide is a strong oxidizing agent; Do not store with oxidizers is a common constraint. Consider contacting your local fire department or the chemistry department of a nearby university for advice on the safe storage of chemicals.
Note; While I'm not endorsing them, I know that Fisher Scientific provides a range of chemical storage cupboards of this type, and has information on which are good for acids, corrosive chemicals and flammable solvents.

Thank you. I have to look if I can find some inexpensive and not too heavy containers for this purpose. I see Fisher Scientific's are >3000$ for a small one. I don't have that kind of budget at home.

Meh, personally I'm not worried about the vapours. You'll be sane and be doing this in an at least somewhat ventilated space, right? Don't do anything outright stupid, but also don't fear the substances so much that your defences cause you to trip and break a toe. $\ce{HCl}$ and $\ce{NaOH}$ aren't exactly harmless, but they're not dimethyl mercury either. There are loads of each of those ions in your body, but as another answer mentioned, $\ce{NaOH}$ doesn't play nice with eyes. (It also turns the oils in your skin into soap, but that's generally not a lifetime liability.)
Like others have said, hydrogen peroxide won't like light. It develops oxygen, so storing it in a gas-tight container in conditions where it's likely to decompose, could cause trouble. I'll get back to this later.
You might want to keep your acid undiluted. If you dilute it, etching can become slow.

AFAICT, you want polyethylene containers throughout. I wouldn't go with glass as a dropped glass container doesn't simply harmlessly bounce off the tiled floor like a plastic container does. I guess a polypropylene ice cream box would work as an etching bath too, but IIRC polypropylene doesn't like the acid all that much (but I haven't tried, and it might well be good enough.)
By all accounts I can find, you really, really want to keep the used etchant - it'll become a cupric chloride solution (which you may have to regenerate occasionally either by aerating or adding hydrogen peroxide). That means you have yet another storage problem. This is no longer pure $\ce{H2O2}$, so you have to think about where the evolved oxygen goes. And if you let the evolved oxygen vent, hydrogen chloride will also vent, attacking any metal nearby. The electronics people have tales of ruined tools sharing a garage with etchant solution. I've resigned myself to keeping the etchant indoors (in my kitchen in fact! where the cost of damaged stuff is least) which is at least better-ventilated than my garage.
And whatever you do, don't pour the beautiful blue cupric chloride down the drain. It'll eat your sink and bring forth chemical armageddon. Don't water your garden with it either, as it's a herbicide. I guess that leaves the options of making it someone else's problem (pay someone to dispose of it for you), or use it to electroplate something with copper, using an inert (not just more copper) anode. Plate faster than the solution etches.

For handling $\ce{NaOH}$ always use goggles (I know myself a case of a student that lost some 15 years ago her vision in the accident with $\ce{NaOH}$. In case of acid you'll get burn but with a bit of luck it will pass). In the case of even a diluted solution of a strong base, your eyes are gone.
The other thing: check your gloves for the chemicals the cover. The usual latex gloves cannot protect against bases perfectly; nitrile gloves (usually bright blue) are much better in this case.
Karen Wetterhahn had gloves on her hands when she worked with dimethylmercury solution. However, she didn't know that this compound goes through a latex in the blink of an eye. Of course, none of the compounds you have mentioned are that dangerous, but it's your health.

Thank you for your precious suggestions. Exactly what kind of goggles? Yeah I thought nitrile gloves are better (found a pack of 100 on farnell for a good price).

Different labs I've been in use different types of goggles. Standard lab safety glasses are fine for $5 by most people's count, but for some things I do prefer my (less comfortable) $15 goggles that fully seal my eyes in.

@user51166: (Standard lab safety glasses are made from polycarbonate, just so you know what to look for.)

Many people prefer nitrile gloves over latex because the latter can lead to hypersensitivity and even allergies.

Storage considerations for chemicals are often noted on their Material Safety Data Sheet, or MSDS. All suppliers that I know of include a MSDS in the packaging when they ship chemicals. You can find lots of good sources of free MSDSs online through Google. Just google "Hydrochloric acid MSDS". I would look for one that appears to come from a reputable source, such as a commercial supplier.
As an example, I have linked the MSDS for concentrated HCl from Fisher Scientific. Section 7 of this MSDS is titled "Handling and Storage".

Thank you. I though these MSDS were quite "universal" (I saw on Google there was a large site containing all documentation about dangerous materials) - not specific to each fabricator. Would sodium persulfate be safer to store than hydrochloric acid by the way ?

Each manufacturer is required to provide them by law. They will almost always have the same information; I just use Sigma-Aldrich's sheets.


text:
Are there any general chemical rules for choosing initiators for a polymerization reaction?
I'm working with some very basic free-radical initiated polymerization reactions. For the most part, I just follow a recipe I'm given (by a member of my group or the broader literature), it works, and I don't worry too much about the chemistry of why it works. Clearly, that will only work for so long, and I'd like to better understand the chemistry of what I'm doing.

So what general rules can I use to predict what kind of initiator will polymerize a given monomer?

To give a concrete example of what I mean: I know that some persulfate-based initiators will polymerize some acrylamide-based monomers. Can I generalize that to all persulfate initiators and all acrylamide monomers? What about methacrylamides? What about styrenes? Essentially, at what point have I changed the reactants so much that the chemistry doesn't work anymore?
Odian's Principles of Polymerization (4th ed) reviews some things that have been done, but doesn't discuss any general rules. Are there other good sources for this information?
EDIT: I am looking specifically for chemical reasons. Are some initiators better at reacting with some monomers? Are there any issues like side reactions I should take into account?
EDIT(2): from @Ben Norris's answer and the ensuing comments, it sounds like, as long as I can get the initiator to radicalize, and everything is soluble in the solvent, then I will be able to use any initiator with any monomer. Is that correct? I've found another textbook on this, Polymer Chemistry: An Introduction by Malcom Stevens, and that seems to support this.

It's been my experience that initiators are chosen for: 1) initiation temperature, 2) solubility, and 3) cost.
1) Initiation temperature - Most radical initiators are stable at room temperature; they need to be heated to dissociate. For example, the half-life of benzoyl peroxide at room temperature is very long, but it is about 30 minutes at 92 °C and 1 minute at 131 °C. However, some monomers boil or decompose at temperatures lower than those temperatures. AIBN (azobisisobutyronitrile) decomposes above 65 °C, so less heat is needed. Both benzoyl peroxide and AIBN are suitable for bulk polymerization of most monomers, including acrylamides, acrylates, styrenes, etc.
2) Solubility - Benzoyl peroxide and AIBN are soluble in organic solvents, which include most monomers for bulk polymerizations. However, if you want to perform your polymerization in water, for example if you are doing an emulsion polymerization, then you need a water-soluble initiator, like a peroxodisulfate salt.
3) Cost - Every initiator has a price, even if you are only using tiny amounts. AIBN may have a lower activation temperature than benzoyl peroxide, but AIBN is also more expensive. There are other more specialized inhibitors that are even more expensive.

This is very helpful. I've clarified that I'm more interested in chemistry-specific reasons. Are some initiators better at reacting with some monomers? Are some imitators particularly prone to certain side reactions?

Initiators are usually present in such low concentrations that initiator induced side reactions are not serious problems. I'm not sure about incompatibility. All of the initiators I listed work with the monomers you listed. For more information, my reference book of choice is the Handbook of Polymer Synthesis from Wiley.

What Ben Norris wrote, is first "chemistry-specific", and second all what can be said.

Ben and @Georg, I've added a question to clarify this, but this sounds like any initiator will work with any monomer, as long as they are all soluble and I can get the initiator to produce radicals. Is that correct?

@ColinMcFaul - That's about the size of it.

Apart from cost, solubility and temperature, the only question is wether the initiator is able to initiate the chain reaction. Because afterwards, it doesn't matter any more, the polymeric radical reacts with more monomer regardless of what the original initiator was.
So the first question is whether your polymer radical is reactive enough to work, but not too reactive, so it would make all kinds of radical (side) reactions instead of just extending the chain.
So, except in the latter case where your monomer was bad at radical polymerisation anyway, the only initiator that wouldn't work is one making very stable radicals. I don't think anyone would advertise those.
There might be a few corner cases, where the addition of the first molecule of a specific monomer takes too long, so the small and highly diffusive initiator radicals have a good chance to recombine before starting a new chain. I don't know of any actual examples for that.


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How does a Brunauer–Emmett–Teller device function?
The Brunauer–Emmett–Teller (BET) theory is comprehensible.
My question is more geared towards what is physically and chemically happening within a BET device to get a reading?  How does a BET device actually go about measuring the physisorption or chemisorption of a material?  Why does this process tend to take multiple hours?

There are no such things such as 'BET devices'. What you can find are gas (or vapor) adsorption apparatus which measure adsorbed amounts by volumetric or gravimetric methods.

BET analysis on zeolithes represents a typical example.  Dependent on the size of the particles you have, you select a glass cell of about a centimetre; for a powder a smaller cell, for more coarse pellets you don't want to grind a bigger one.  You need to outgas the samples first, i.e. by applying vacuum and elevated temperatures already adsorbed gases and moisture is removed from the surface of the material to be analyzed as well as the inner walls of the sample cell. This already may take time to be accomplish -- because you do not want to destroy or alter the material in question by too harsh conditions.
The actual measurement requires the sample cell being kept at a constant temperature.  An initial determination takes place while the sample cell is immersed in a Dewar of liquid nitrogen.  This low temperature provides better conditions for the adsorption experiment.
A first blank is measured with Helium gas, entered via a graduated piston. Helium typically does not adsorb on surfaces (at this temperature).  This allows to determine the dead volume of the measurement cell.  In the second step, Nitrogen gas enters the measurement cell, again in a controlled manner (graduated piston, pressure control).  Again, this takes time as you want to establish the equilibrium between the adsorbent and the amount of gas adsorbed in function of gas pressure and sample temperature.
A controlled entry of Helium gas is necessary as a second determination of the dead volume to complete the calibration eventually leading to the read-out for our substance at given temperature.
Of course, dependent on what kind of isotherme (BET/Langmuir) is to be determined, Nitrogen gas as a load may be replaced by other gases, like methane.  Beside determination of the surface area of a solid, the adsorption properties may be determined in a series of more elevated temperatures, too.


text:
Do hydrogen halides (HX) form strong or weak acids?
I know that when dissolved in water, hydrogen halides ($\ce{H-X}$), become acids. I know that $\ce{HF}$ forms a weak acid, but can it be generalized that all $\ce{HX}$ acids do so?

No, it can’t be generalised. Amongst the haloacids, only $\ce{HF}$ is weak.
Acid strength can be determined by the stability of the conjugate base (for most acids, the conjugate base is what you get when you remove $\ce{H+}$ from the acid). In this case the conjugate bases are the halides.
As the size of the halide ion increases, the charge is dispersed over a larger volume, and thus the ion is more stable. In this manner, one can say that the order of acid strengths is $\ce{HF<HCl<HBr<HI}$ (I’m not considering the radioactive ones here; they fit in the same trend though).
$\ce{HCl, HBr, HI}$ are all strong acids. $\ce{HCl}$ has a $\mathrm{p}K_\text{a}$ of −7 (the smaller the $\mathrm{p}K_\text{a}$ is, the more acidic it is), $\ce{HBr}$ is approximately −9, and $\ce{HI}$ is also approx −9.
On the other hand, $\ce{HF}$ is 3.2. The $\ce{H-F}$ bond is pretty strong, so it has a smaller tendency to dissociate—making it a weak acid (not to say that it isn’t dangerous).

Ah, of course about HCl and everything. I originally asked if they were all weak, then remembered about HF and forgot about the others. Thanks for relating it with other things like pKa.

I'd have said $\ce{HF \ll HCl<HBr<HI}$ with regards to the acidity myself. :) As another example of the tenacity of hydrogen bonding in $\ce{HF}$, consider ammonium bifluoride...

No: HCl is a very strong acid, and HBr even more so. HF is the only weak acid. This is because it binds hydrogen much more strongly then the other halides.


text:
The mole is used extensively in chemistry, why not elsewhere?
In chemistry, the unit "mole" is used in many places, and is an important tool when working with quantities.
I have not seen it used elsewhere--for example, we never count people or TVs with the unit.
What makes the mole so useful in chemistry, and why isn't it used elsewhere?

Because the mole is not a convenient basis for people and TV sets (for instance, one mole of people would be approximately 86 trillion (short-scale) times the population of earth, or put another way the earth contains a human population concentration of approximately 12 femtomoles/planet), and because there's no connection back to atomic mass, which is what makes the mole useful in the first place.
This was not part of your original question, but addresses Manishearth's modifications:
Because atoms and molecules react in discrete units (cf. the law of definite proportions), the number of particles is more fundamental than the masses of those particles. The mole is useful in chemistry because it is defined such that the mass of Avogadro's number of particles (atoms or molecules) of a substance (the molar mass) is equal to the sum of the numerical values of the atomic masses of its constituent elements. The value of the mole is calibrated such that the molar mass of carbon-12 is exactly 12 grams per mole. On the periodic table you will note that the atomic masses given are non-integral (often markedly so, e.g. chlorine at 35.45 g/mol). This is because the atomic masses represent a weighted average of isotopic masses, and because the proton and neutron masses aren't quite equal.
The mole makes stoichiometric calculations particularly convenient as we can measure out, for instance, 16 grams (one mole) of $\ce{^{16}O}$ and 2 grams (two moles) of $\ce{^1H}$ and obtain one mole of $\ce{H2O}$. In the case of water, oxygen makes up 89% of the molecule by mass, but is only one of the three atoms present in the molecule. The mole is necessary to transpose stoichiometries involving single atoms and molecules (with their tiny masses) into the everyday world of grams and litres.

Yep, would've been a shame to see this go to waste.. :)

$8.6*10^{13}$ more than earth's population?! Heh, it humors me that I always thought Avagadro's as so big as to be incomparable to anything. Quantifying with 86 trillion times earth's population is suddenly even more absurd. Thanks for the chuckles. +1

@ChrisK - Well, there are about 4.9 octillion IPv6 addresses for every human being, which makes Avogadro's number look positively miniscule.

A mole of stars would be all the stars in all the galaxies we can see, and more.

To satisfy my curiosity, I calculated the length for a mole of dimes in a stack. It would just about be the diameter of the Milky Way.

Out of interest, I tried to come up with a natural use of 1 mol (such as "1 mol of x is about 5 of (another thing)" that doesn't involve atoms/molecules at all. The closest thing I could find is "the number of bacteria on earth is about $5*10^{30}$, or about ten million moles. But that's still many orders of magnitude off!
Source: http://news.bbc.co.uk/1/hi/sci/tech/158203.stm

The Si-unit "mol" is used in chemistry for three reasons:

The quantity of atoms/molecules is the logical unit of chemistry (Reactions have to be balanced).
Atoms/molecules can't be counted (It works for people and TV's), but have to be measured using mass or volume, and then converted into a quantity (e.g. mol/kg)
The quantity of atoms/molecules involved in a typical experiment or measurement is very large, making a shorthand useful (Compare this to the Astronomical Unit which is a shorthand for a large number of km).

The mol (or a similar unit) isn't used elsewhere because:
Anything larger than chemicals can be measured with whichever measure seems to be the most logical. Chemistry is unique in requiring chemical quantities to make sense of the world but at the same time these quantities can not be determined directly.

Apparently mol is huge unit and that makes in not very useful for normal life, however there is another simple reason: there is no balance or cylinder that can measure mol. The unit has its logic in sense of reactions (equimolar etc.) but still every single experimental chemist has to start with weight or volume when he prepares his synthesis.

The mole is useful in chemistry because it relates two popular but inconsistent units of mass - the atomic mass unit (u) and the gram. One mole of atoms of an element with atomic mass x u, has a mass of x grams.
We don't use atomic mass units outside chemistry, so we also have no need for the mole.
This of course raises another question "why do chemists use atomic mass units?". Perhaps it's because they measure quantity in moles :P Really I think it's just an accident of history from the days when people thought atomic masses were all integer multiples of the mass of hydrogen.


text:
Does borax do anything more for "boosting" detergent than adding active oxygen would?
Borax, $\ce{Na2B4O7}$, is often marketed as a "laundry booster" under the brand "20 Mule Team Borax".  The unit crystal of borax can be seen below.

Other laundry products in the past have added sodium perborate (which can be produced from borax, hydrogen peroxide, and sodium hydroxide) as a bleaching additive.  The "peroxide" portion can readily be seen below.

Through what mechanism does the plain borax "boost" the detergent?  Would it be more effective to throw some weak peroxide and a small amount of base into the laundry to form the perborate instead of using plain borax?

Borax was used as a "builder" in washing formulations long before perborate was used as a bleach. "Builder" is the terminology for additives which reduce the influence of water hardness. Other examples of builders include trisodium phosphate, soda, polyphosphates (like Graham's salt) or today most often aluminosilicate ion exchangers (Phosphates are more or less forbidden nowadays because of eutrophication of rivers and seas).


text:
Are all cubic point groups encountered?
My symmetry lecture notes state that there are seven high-symmetry 3D point groups, which have more than one rotation axis of order greater than 2: $T$, $T_d$, $T_h$, $O$, $O_h$, $I$, $I_h$. I sometimes see them called the “cubic point groups”.
My point is: most molecules (or molecular fragments) I have encoutered in these groups are always $T_d$ (tetrahedral) or $O_h$ (octahedral). What are examples of molecules in the other groups?

Just a quick correction: I would only call the tetrahedral and octahedral point groups cubic (which means they're compatible with a cubic crystal lattice) - the icosahedral point groups are not.

have a look at  http://molecule-viewer.com where there are examples of molecules from all point groups. However, its designed test your ability to work out point groups not just look up examples.

I refer you to the second link in my answer for How does one recognized Td/Oh symmetry in molecules?, which (according to the abstract) describes some high symmetry species of unusual point groups $T$, $O$ and $I$. Additionally, the molecule [6.6]chiralane has $T$ symmetry.
Buckminsterfullerene is a well known $I_{h}$ symmetric molecule.
Don't know of any $T_h$ molecules but I'm confident they can be devised.

At least six of these seven point groups been encountered. You can see examples on the Otterbein symmetry gallery. The one that's missing from this gallery is $I$, although as Richard pointed out, people have theoretically predicted some molecules with this symmetry to be stable.


text:
Can class D fire extinguishers be used for other fire types than metal fires?
Burning metals like sodium should not be extinguished with standard fire extinguishers, but they require either sand or a special class D fire extinguisher.
Now, if there is a fire in the lab and the closest fire extinguisher is one of those class D extinguishers meant for metal fires, can those be used for other types of fires?
The main case where there might not be enough time to get specific equipment is probably when people are on fire. Can a class D extinguisher be used in this case, if it's the closest available one?

Tricky question. There's an important distinction between extinguishers that are safe and effective on non-prescribed fire classes, and extinguishers that are safe to use on living human beings. Many of the Class D formulations sound like they could be highly irritating if they were to come into contact with eyes or mucous membranes (powdered graphite, copper, salts).

That said, getting sprayed with powdered salt is probably preferable to being on fire (unless you're a slug or snail), however that sort of compromise is indicative of a lab safety failure.

If a person is on fire, you should use the safety shower in your lab to extinguish the flames. That is a basic safety rule that should be covered in any lab safety class.

@scientifics Safety showers aren't the only equipment in a lab that can be used to extinguish people on fire, fire blankets and fire extinguishers can be used as well. I would usually choose the shower if I have a choice, my question was more targeted towards situations where that isn't possible for some reason.

@scientifics I would definitely not go into the safety shower to extinguish fire on me if one of the following was true: 1) the safety shower is more than 10 feet away, 2) there are significant obstacles between me and it (stools, litter, people, boxes, etc.), or 3) the fire was stared by a pyrophoric or water reactive substance. In those cases, smothering by the good-old stop-drop-roll method is superior.

Unfortunately, as far as I can tell, specific class D fire extinguishers aren't even suitable for all types of class D fire, much less other types of fire.
The copper powder variety in particular has limited effectiveness against fires that aren't lithium-based, and the thermoplastic additives in the sodium chloride-based extinguishers may also catch fire.
Dry sodium carbonate-based class D extinguishers should be generally effective but suboptimal against many types of fire, but I'd say in general, even if you have to go a little further, you should go for the proper sort of extinguisher to minimise accidental horrible disasters.

This answer is the most dangerous nonsense I have until this day read on SE. The most important consideration in fire fighting is SPEED! Empty the slightly less suitable extintuisher on the fire (they usually last 15 seconds!), THEN fetch something else to  make sure the fire doesn't start and spread again. The ONLY extinguisher that really doesn't work is an ABC type against a metal fire.

@Karl please consider writing that an answer, so that it can be voted upon as well - if you explain what you think is wrong with the accepted answer then it might help convince people, too. Your comment already serves as a warning and generally we don't remove answers simply for being wrong, we downvote them.


text:
How do you determine the amount of acrylamide in foods such as baked fries?
There are many food studies which determine the concentration of a substance in a food item.
For example, this study determines the concentration of acrylamide in solid foods like fries.
By what experimental technique is this done?

Check out the methodology page for the FDA section on acrylamide quantitation, which goes into considerable detail.
In short, they start by mashing up some fries, and then extract and purify a fraction containing the acrylamide through a number of steps leading up to liquid chromatography. They then use quantitative mass spectrometry to determine the amount of acrylamide present.
They add a small known concentration of isotopically labelled acrylamide (i.e. acrylamide with all the carbons as carbon-13, which has very low natural abundance and results in a molecule that is approximately 3 g/mol heavier). This shows up in the mass spectrum as an offset peak. Then they can integrate the normal acrylamide mass spectrum peak and the labelled peak and obtain the concentration from the relative peak areas.


text:
Oxidative chlorination mechanism (sulfide to sulfonyl chloride)
I cannot find much guidance in the literature. Any ideas as to what the mechanism of this reaction might be?

The mechanisms of oxidations of this type are often difficult to study since there are multiple pathways they can follow. For instance, I can draw valid-looking dipolar and radical mechanisms for this reaction. When I'm back in my office and have access to ChemDraw, I will post both as answers.

Regarding the species involved, $\ce{Cl_2}$ in water should disproportionate to $\ce{ClO^-}$, and possibly also create $\ce{CH_3COCl}$. If you think about this species as well, what reaction you can draw?

As Ben Norris already pointed out, these processes are typical a bit complex. It is however conceivable that the direct oxidation of the thioether 1 is among the first steps involved.

While the oxidative cleavage of PMB (= 4-methoxybenzyl) ethers, often applied as protecting groups, is assumed to take place in the methoxy-substituted arene, here the electron transfer mostly likely proceeds from the sulfur atom, leading to the sulfur-centred radical cation 2.
In the presence of a nucleophile, this species is likely to break down en route to the radical 3.

Whether a disulfide, formed upon dimerization of 3, will take part in the further course of the reaction is beyond these musings.


text:
Why are many drugs in the hydrochloride form?

Possible Duplicate:
Why are solid phase organics often conjugated to HCl?

I'm wondering why a lot of drugs are in this hydrochloride form.
E.g.,

Cetirizine hydrochloride
Methylphenidate hydrochloride
Tramadol hydrochloride

I'm really curious to know why this form is so common for the presentation of a drug.


text:
Negative retention for chloride in hydrolysate filtration
I tried to fractionate proteins as a function of their size with an ultrafiltration membrane. I observed a higher concentration of chloride in my permeate than in my feed or retentate. The proteins do not pass through the membrane and the filtration has been done at pH = 7. I asked, how is it possible to have a higher concentration in the permeate than in the feed or retentate?
I found this explanation but I'm not sure about it :

It may be explained by the Donnan effect. Negatively charged proteins are retained in the retentate. To compensate the positive charges from the cations that are on the other side of the membrane, they attract the chloride.

Is it a possible way to explain this issue?

It depends. That effect relies on both sides of the membrane being immersed in a liquid phase. Is that the case in your setup, or are you using the membranes in a traditional filter setup with one side dry/not immersed?

A combination of osmosis and Gibbs-Donnan effect is the correct explanation.
With a semi-permeable membrane such as the one you are using, or a cell membrane, or in dialysis, water can move freely as can chloride and other small ions, however large molecules such as proteins cannot.
A simplified explanation is this: in osmosis, the tendency is for the osmotic pressure and charge to equalise on either side of the membrane. Proteins usually have a negative charge and exert a large osmotic pressure. Thus, to balance this out, small anions will diffuse to the opposite side to where the proteins are. Water will diffuse until the osmolality is equal on either side of the membrane unless there is a hydrostatic pressure forcing it to the other side. In the absence of hydrostatic pressure, you will end up with a higher concentration chloride solution on one side and a higher concentration protein solution on the other side.
Here is a diagram:

If you assume that the small particles you can see are chloride and the big particles are your proteins, you can see that the chloride concentration in the permeate is higher than the retentate and higher even than the feed.


text:
How can one optimise the formation of copper acetate?
Mixing solid copper, $5\%$ vinegar, and hydrogen peroxide ($\ce{H2O2}$) causes copper acetate to form. The process will occur very slowly without hydrogen peroxide. Adding $\ce{H2O2}$ speeds the formation of copper acetate.
How can I test the solution to see if all of the acetic acid has been converted into copper acetate?
If possible, provide answers that employ only household items or readily available devices.

Your question and title do not correspond, IMHO. I'll answer one of both:
Vinegar is mostly water, with some acetic acid ($5\%$). A classic experiment, performed by a lot of kids at school, is to mix vinegar with baking soda. The baking soda neutralizes the acid and produces carbon dioxide. The carbon dioxide is what you're looking for: it's the sign that the reaction is taking place.
Pour a pinch of baking soda inside your solution. If it makes bubbles on contact, your solution still contains a lot of vinegar. If you don't see any bubbles, even when stirring, most of the acid has been consumed.